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A  MANUAL  OF  CHEMICAL  ANALYSIS 


BY   THE  SAME  AUTHOR 

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With  224  Diagrams.     Crown  8vo,  $  2.00. 

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LONGMANS,   GREEN,    AND   CO. 

NEW  YORK,    LONDON,    BOMBAY,    AND   CALCUTTA. 


A    MANUAL 

OF 

CHEMICAL    ANALYSIS 

QUALITATIVE   AND    QUANTITATIVE 


BY 


G.   S.   NEWTH,   F.I.C.,   F.C.S. 

DEMONSTRATOR   IN   THE   ROYAL  COLLEGE   OF   SCIENCE,    LONDON 
ASSISTANT-EXAMINER   IN   CHEMISTRY,    BOARD   OF   EDUCATION 


UNIVERSITY 

OF 

SlUFORN 


EIGHTH  IMPRESSION 


LONGMANS,    GREEN,    AND    CO, 

91    AND   93    FIFTH   AVENUE,    NEW   YORK 

LONDON,  BOMBAY,  AND  CALCUTTA 

1909 


Engineering 
Library 


AW         •  *~* 


P 


PREFACE 

A  MAN  once  brought  his  son  to  the  Royal  School  of  Mines — now 
the  Royal  College  of  Science— with  the  request  that  he  might  be 
taught  to  "  do  copper."  He  did  not  want  his  boy  to  "  waste  his 
time  learning  about  oxygen  and  hydrogen,  and  all  that"  but  he 
wished  him  simply  to  learn  to  "do  copper." 

Although  seldom  expressed  with  such  refreshing  candour,  the 
desire  to  do  analysis  without  learning  more  than  the  minimum 
amount  of  chemistry  is  still  very  prevalent  ;  and,  unfortunately, 
chemical  analysis  is  a  subject  which  may  be,  and  frequently  is, 
taught  and  practised  in  such  a  manner  as  to  degrade  it  to  the  level 
of  a  purely  mechanical  and  often  quite  unintelligible  series  of  rule- 
of-thumb  operations. 

I  hope  that  the  student  whose  aspirations  rise  no  higher  than  to 
learn  to  do  analysis  in  this  fashion,  will  not  find  this  book  suitable 
for  him.  I  have  done  my  best  to  make  it  as  little  of  a  cram-book 
as  possible,  but  have  endeavoured  to  teach  analytical  chemistry  as 
well  as  analysis — that  is,  the  theoretical  as  well  as  the  practical  side 
of  the  subject. 

With  this  object  in  view,  I  have  carefully  avoided  the  use  of  those 
symbolic  abbreviated  expressions  (slang  formula,  they  might  be 
termed),  such  as  H2O  (oxalic  acid),  H2T  (tartaric  acid),  HA  (acetic 
acid),  etc.,  which  are  becoming  so  common,  and  which,  so  far  as 
the  student  is  concerned,  foster  those  very  evils  of  cramming  which 
we  as  teachers  are  striving  to  combat.  _  If  such  symbols  as  these 
are  permitted  and  recognised,  why  not  H2S,  HN,  H3P,  for  sulphuric, 
nitric,  and  phosphoric  acids  respectively  ?  And  then,  perhaps, 
some  such  hieroglyphic  as  H2S,  HN,  for  sulphurous  and  nitrous 
acids. 

For  the  busy  chemist  to  make  use  of  such  shorthand  signs  in 
the  privacy  of  his  own  note-books  is  one  thing,  but  to  print  them  in 

a  2. 

on  t  ^a 1 


vi  Preface. 

a  text-book  intended  for  the  learner,  and  thereby  put  them  on  the 
same  footing  as  such  chemical  formulas  as  H2SO4,  H2C2O4,  etc.,  is 
quite  another  thing,  and  is  open  to  the  taunt  of  chemistry  made  easy. 

The  importance  to  the  student  of  making  careful  notes  of  his 
analytical  work  while  in  progress,  cannot  well  be  overrated,  as, 
perhaps  more  than  anything  else,  this  is  calculated  to  develop  in 
him  those  habits  of  exact  observation  which  are  essential  qualities 
in  a  scientific  man.  While  insisting  on  this  point  more  than  once, 
I  have  purposely  refrained  from  giving  anything  of  the  nature  of  a 
specimen  of  notes,  because  notes  which  are  made  according  to  a 
stereotyped  pattern  are  practically  of  no  value.  Unfortunately,  it 
is  the  custom  in  most  laboratories — perhaps  a  necessary  custom, 
but  none  the  less  unfortunate — to  require  from  the  student  some 
written  account  of  his  analysis  by  way  of  proof  that  he  has  con- 
scientiously done  the  work.  The  result  of  this  is  that  the  notes 
which  he  makes  are  taken  primarily  with  a  view  to  furnish  this 
required  evidence.  Points  which  ought  to  be  observed  and  noted, 
points  which  he  does  not  at  the  time  understand,  are  passed  over, 
while  others  which  are  familiar,  and  which  are  at  once  recognised 
as  useful  evidence,  are  written  down.  Indeed,  as  every  teacher 
has  experienced,  the  evidence  is  often  more  or  less  fabricated  after 
the  completion  of  the  analysis. 

The  consequence  of  this  method  is  the  stereotyped  product  so 
painfully  familiar  to  all  who  are  engaged  in  teaching  or  examining, 
namely,  a  certain  number  of  regulation  tests,  arranged  with  a 
semblance  of  method  in  the  everlasting  three  columns,  headed 
Experiment,  Observation,  Inference,  followed  by  a  more  or  less 
slovenly  copy  of  the  time-honoured  tables. 

The  valuelessness  of  such  a  written-out  record  must  be  patent 
to  all  teachers,  and  yet  we  still  continue  to  accept  it  in  lieu  of  a  few 
real  notes,  less  quickly  and  easily  examined  no  doubt,  because  less 
stereotyped,  but  of  infinitely  more  value  to  the  student  himself. 

In  order  to  gain  as  much  space  as  possible  for  purely  analytical 
matter,  and  still  to  be  able  to  include  within  the  limits  of  a  con- 
venient volume,  both  qualitative  and  quantitative  analysis,  I  have 
carefully  excluded  all  merely  descriptive  details  which  have  no 
direct  bearing  upon  analysis.  For  example,  none  of  the  properties, 
either  chemical  or  physical,  of  such  metals  as  those  of  the  alkalies 
or  alkaline  earths,  are  made  use  of  in  ordinary  analytical  processes 
for  the  detection  or  recognition  of  these  elements,  therefore  any 


Preface.  vii 

descriptive  account  of  their  properties  in  the  elemental  state  is 
entirely  out  of  place  in  a  book  which  is  intended  only  to  teach 
analytical  as  distinguished  from  general  chemistry. 

In  the  second  portion  of  the  book,  devoted  to  quantitative 
analysis,  I  have  confined  myself  to  a  comparatively  small  number 
of  well-tried  typical  methods  and  processes,  preferring  to  describe 
and  explain  in  tolerably  full  detail,  a  few  quantitative  determinations 
in  each  of  the  various  sections,  such  as  shall  furnish  a  thoroughly 
sound  course  of  practical  study,  rather  than  to  attempt  to  cover — 
necessarily  in  more  sketchy  outlines — a  wider  range  of  subjects. 
Whether  the  choice  of  processes  and  examples  I  have  made  is  as 
good  or  as  representative  as  it  might  be,  will  no  doubt  be  a  matter 
for  difference  of  opinion,  and  as  the  book  has  not  been  written  to 
meet  any  particular  syllabus,  this  is  probably  a  point  upon  which 
hardly  any  two  teachers  will  exactly  agree. 

The  illustrations  throughout  have  been  made  from  original 
photographs  of  the  actual  apparatus  employed  in  the  various 
analytical  operations  described. 

G.  S.  N. 

ROYAL  COLLEGE  OF  SCIENCE,  LONDON, 
June,  1898. 


TABLE    OF   CONTENTS 

BOOK  I. 

QUALITATIVE  ANALYSIS. 

PAGE 

CHAPTER  I.     Preliminary  exercises — Filtration;  Solution;  Evapo- 
ration ;  F^ision  ;  Precipitation  ;  Ignition  ;  Neiitral- 

isation I 

,,       II.     Analytical  classification 13 

,,      III.     Reactions  of  the  metals  of  Group  V 19 

APPENDIX  TO  CHAPTER  III.     The  rare  metals  of  Group  V.    .  26 

CHAPTER  IV.     Reactions  of  the  metals  of  Group  IV 30 

V.                   „                   „                   „                   IIlA 36 

VI.                    „                    „                    „                    HlB 51 

,,        VII.     The  phosphates 63 

APPENDIX  TO  CHAPTER  VII.     The  rare  metals  of  Group  III.  70 

CHAPTER  VIII.     Reactions  of  the  metals  of  Group  II.,  Division  I  75 

IX.            „                  ,,                 ,,             II.,  Division  2  89 

APPENDIX  TO  CHAPTER  IX.     The  rare  metals  of  Group  II.    .  108 

CHAPTER  X.     Reactions  of  the  metals  of  Group  1 112 

APPENDIX  TO  CHAPTER  X.     The  rare  metals  of  Group  I.  .     .  118 

CHAPTER  XI.     The  non-metals  and  their  acids — The  halogens  .     .  120 

„        XII.     Sulphur 137 

,,      XIII.     Nitrogen  and  phosphorus 144 

,,       XIV.     Carbon,  silicon,  boron 153 

,,         XV.     Systematic  detection  of  the  acids 171 

,,       XVI.     Preliminary  tests  and  operations  in  a  systematic 

analysis  j  general  analytical  tables        ....  178 

,,     XVII.    The  results  of  a  qualitative  analysis IJ56 

BOOK  II. 

QUANTITATIVE  ANALYSIS. 

PART  I.    GRAVIMETRIC  METHODS. 

SECTION  I.     Preliminary  manipulations 189 

(i)  Weighing. 


Table  of  Contents. 


(2)  Drying  and  weighing  a  filter. 

(3)  Estimation  of  water  of  crystallisation. 

(4)  Determination  of  the  ash  of  a  filter. 

(5)  Preparation  of  pure  salts. 

SECTION  II.     Typical  gravimetric  determinations  of  metals       .     221 
Al,  Cr,   Fe,  Ca,  Mg,  Cu,  Ag,  Pb,  Zn,  Mn,  Ni,  Co,  K, 

NH4,    Sn,  As,  Sb,  Cd,  Hg. 

SECTION  III.     Typical  gravimetric  determinations  of  acids  .      .     258 
H2S04,  HC1,  HBr,  CO2,  HNO3,  H3PO4,  SiO2. 

S-ECTION  IV.     Exercises  in  gravimetric  analysis 265 

Silver  coin,  solder,  German  silver,  bronze,  dolomite,  zinc 

blende,  silicate. 
SECTION  V.     Electrolytic  methods  of  analysis 292 

PART  II.    VOLUMETRIC  METHODS. 

SECTION  I.     Preliminary  manipulations 300 

(1)  Instruments    for    measuring    liquids  —  Calibration    of 
vessels. 

(2)  Standard  solutions. 

(3)  Methods  for  ascertaining  the  completion  of  volumetric 
reactions — Indicators. 

SECTION  II.     Volumetric    methods  based    upon    saturation- 
Alkalimetry,  acidimetry      3l6 

Typical  analyses  by  means  of  standard  acids  and  alkalies. 
SECTION  III.     Methods  based  upon  oxidation  and  reduction     .     331 
Potassium  permanganate — Titration  with  iron,  ferrous  salt, 

oxalic  acid. 
Typical  analyses  by  means  of  permanganate. 

(1)  Iron  ores. 

(2)  Available  oxygen  in  manganese  ores. 

(3)  Estimation  of  tin. 

(4)  Estimation  of  calcium. 

Potassium  dichromate — Titration  \v\\hferrous  salts. 
Analyses  by  means  of  dichromate. 

(1)  Estimation  of  iron  in  iron  ores. 

(2)  Estimation  of  chromium  in  chrome  iron  ore. 
Iodine  solution— Titration  with  thiosulphate  and  arsenious 

oxide. 
Estimations  by  means  of  iodine. 

(1)  Antimony. 

(2)  Arsenic. 
(2)  Tin. 

Estimations  by  means  of  iodine  and  thiosulphate. 


Table  of  Contents.  xi 

PAGE 

(1)  Sulphur  dioxide  in  a  sulphite. 

(2)  Available  chlorine  in  bleaching  powder. 

(3)  Estimation  of  manganese  dioxide. 

SECTION  IV.     Methods  based  upon  precipitation      ....     360 
(I.)  Precipitation  with  silver  nitrate — Chlorine,  cyanogen. 
Indirect    estimation  .of    nitric    acid,    calcium,   carbon 

dioxide,  and  cadmium. 

(II.)  Precipitation  with  sodium  chloride — Silver. 
(III.)  Precipitation  with  ammonium  thiocyanate. 

(1)  Silver. 

(2)  Indirect  estimation  of  chlorine. 

(IV.)  Precipitation  by  means  of  uranium  acetate. 

Estimation  of  phosphoric  acid  in  bone  ash. 
(V.)  Precipitation  with  sodium  sulphide. 

Estimation  of  zinc  in  zinc  ores. 

(VI.)  Clark's  method  for  estimating  hardness  of  water, 
APPENDIX  TO  SECTION  IV.    Estimation  of  copper  by  potassium 

cyanide 374 

SECTION  V.  Gas  analysis 377 

(1)  Estimation  by  absorption  and  subsequent  titration. 

(a) -Carbon  dioxide  in  air  (Pettenkofer's  method). 
(b)  Sulphur  dioxide  in  furnace  gas. 

(2)  Estimation    by  absorption,   and    measurement   of    the 
residual  gas. 

Measurement  of  gases. 

Calibration  of  gas-burettes. 

Correction  of  gaseous  volumes. 

Collection  of  samples  for  analysis. 

Absorption  in  Hempel's  gas-pipettes. 

Analysis  of  mixture  containing  CO2,  CO,  O,  N. 

(3)  Estimation   by   combustion,  and   determination  of  the 
products. 

Hydrogen  by  combustion  with  palladium. 

Hydrogen  by  explosion. 

Marsh-gas  by  explosion. 

Analysis  of  mixtures  containing  H,  CH4,  N. 

The  nitrometer. 

PART  III. 

SECTION    I.       Estimation    of    carbon,     hydrogen,     nitrogen, 

chlorine,  sulphur,  and  phosphorus  in  organic  compounds  .      .     414. 
SECTION  II.     Miscellaneous  physico-chemical  determinations  .     434 
Specific  gravity  of  solids  and  liquids. 
Boiling-point. 


xii  Table  of  Contents. 


Melting-point. 

Vapour  density. 

Molecular  weight  by  freezing-point. 

Molecular  weight  by  boiling-point. 

APPENDIX — 

Table  of  elements  and  atomic  weights 453 

Reagents — System  of  standard  strengths 454 

Table  of  hardness .      .      .      .  459 

Expansion  of  water  between  o°  and  25° 460 

Tension  of  aqueous  vapour  from  5°  to  25° 461 

Table  of  factors  for  the  correction  of  gaseous  volumes    .      .  462 

Table  of  logarithms 466 


CHEMICAL  ANALYSIS 

BOOK    I. 

QUALITATIVE  ANALYSIS. 

CHAPTER   I. 
PRELIMINARY  EXERCISES. 

THE  first  step  that  the  student  must  take  in  approaching  the 
subject  of  analytical  chemistry,  is  that  of  making  himself  practically 
familiar  with  certain  simple  operations  or  manipulations  which  he 
will  constantly  be  required  to  carry  out  in  the  course  of  his  work, 
and  upon  the  dexterous  and  cleanly  performance  of  which  much 
of  his  success  as  aa  analyst  will  depend.  If  he  has  not  had  previous 
experience  in  practical  chemistry,  therefore,  he  should  carefully  go 
through  the  following  exercises. 

i.  Filtration. — The  method  by  which  a  liquid  is  separated 
from  any  solid  substance  with  which  it  is  mechanically  mixed,  is 
most  usually  that  of  filtering  the  mixture  through  porous  paper, 
known  as  filter-paper. 

EXERCISE  i. — Fold  a  circular  filter-paper  into  half,  and  then  at 
right  angles  into  half  again.  Open  this  into  a  cone  having  one 
thickness  of  paper  on  one  side  and  three  on  the  other.  This  cone 
is  then  placed  in  a  glass  funnel  of  such  a  size  that  the  glass  will 
project  slightly  above  the  paper.  The  paper  is  then  moistened 
with  distilled  water,  which  should  not  be  poured  out  of  the  funnel 
again,  but  allowed  to  run  through.  After  being  cautiously  pressed 
into  the  glass  funnel,  the  paper  should  fit  close  to  the  glass  all 
round,  leaving  no  air-spaces.  If  this  is  not  the  case,  either  another 
funnel  of  the  right  angle  (60  degrees)  should  be  selected,  or  another 
filter-paper  folded  so  that  the  cone  shall  be  of  the  same  angle  as 
the  funnel.  This  can  be  done  by  making  the  second  fold  of  the 

B 


2  Qualitative  Analysis. 

paper  not  quite  at  right  angles  to  the  first.  In  this  way  a  cone 
will  be  formed  having  either  a  more  acute  or  more  obtuse  angle 
than  60  degrees,  as  the  paper  is  opened  out  one  side  or  the  other. 
The  funnel  is  supported  by  a  metal  or  wooden  stand. 

Now  place  some  diluted  hydrochloric  acid  in  a  small  beaker, 
and  stir  into  it,  by  means  of  a  glass  rod,  a  quantity  of  finely 
powdered  charcoal.  When  thoroughly  mixed,  pour  upon  the  filter. 
When  slowly  pouring  from  a  wide  vessel  like  a  beaker,  there  is 
risk  of  some  of  the  liquid  being  spilt  by  running  down  the  outside 
of  the  vessel,  as  shown  in  Fig.  i.  If  it  be  poured  quickly,  it  is 


FIG. 


FIG.  2. 


likely  to  splash  over  the  funnel.  To  prevent  both  of  these  acci- 
dents, the  liquid  should  be  poured  down  against  a  glass  rod  held 
lightly  against  the  edge  of  the  beaker,  and  in  such  a  position  that 
the  liquid  does  not  strike  at  once  against  the  apex  of  the  paper 
cone  (Fig.  2). 

The  filtrate  (i.e.  the  liquid  which  passes  through  the  filter)  may 
be  received  in  another  beaker,  which  should  be  placed  close  against 
the  stem  of  the  funnel,  so  that  the  liquid  shall  run  down  against 
the  glass.  In  this  way  splashing  is  prevented.  The  filtrate  should 
be  perfectly  clear,  the  whole  of  the  solid  being  retained  on  the 
filter.  When  all  the  liquid  has  passed  through,  the  charcoal  and 
the  filter-paper  are  both  still  soaked  with  the  hydrochloric  acid. 
In  order  to  remove  this,  and  so  to  make  the  separation  of  the  solid 
from  the  liquid  complete,  the  filter  and  its  contents  must  be  washed 


Preliminary  Exercises. 


with  distilled  water.*  This  is  done  by  directing  a  fine  stream  of 
water  from  a  wash-bottle  into  the  funnel,  working  downwards  from 
the  upper  edges  of  the  paper,  and  so  washing  the  charcoal  down 
into  the  apex  of  the  filter  (Fig.  3).  Each  washing  must  be  allowed 
to  drain  right  through  before  more  water  is  used.  This  must  be 
continued  until  the  filtrate  is  entirely  free  from  acid,  which  may 
be  ascertained  by  allowing  one  or  two  drops  of  it  to  fall  upon  a 
piece  of  blue  litmus  paper. 

In  practice,  the  size  of  the  filter  should  bear  a  rational  relation 
to  the  quantity  of  solid  matter  to  be  separated  from  a  liquid.     This 
is  more   especially  impor- 
tant when  the  material  re- 
tained upon  the  filter  has 
to     be    washed.       If    the 
amount   of  solid  is  small, 
the  filter   used   should  be     / 
proportionately  small,  and   / 
the  washing  operation  will 
be  more  quickly  and  effec- 
tually accomplished  than  if 
an    unduly  large    filter  is 
employed. 

2.  Solution. — This 
term  is  applied  both  to  the 
act  of  dissolving  and  to  the        ; 
product   obtained   by  dis-      j 
solving. 

EXERCISE  2.— Place  a 
little  powdered  potassium 
carbonate  in  a  test-tube, 
and  add  a  small  quantity 
of  water.  In  a  few  moments 
the  salt  will  have  entirely 
dissolved.  The  salt  has 
undergone  solution  in  water. 


FIG.  3. 


The  product  is  a  solution  of  potas- 
sium carbonate.  The  water  is  called  the  solvent.  The  process  of 
solution  is  accelerated  by  heating  the  liquid,  and  it  takes  place  more 
quickly  the  more  finely  the  solid  is  powdered. 

Put  a  similar  quantity  of  potassium  carbonate  into  another 
test-tube,  and  add  a  little  dilute  nitric  acid.  The  salt  again  under- 
goes solution,  the  acid  here  being  the  solvent.  But  in  this  case 
there  is  a  radical  difference.  First,  a  visible  difference,  in  that  the 

*  In  the  following  exercises,  and  in  all  analytical  operations,  distilled 
water  must  always  be  employed;  and  when  beakers,  test-tubes,  etc.,  are 
washed  up  after  use,  they  must  be  finally  rinsed  with  distilled  water. 


Qualitative  A  nalysis. 


act  of  solution  is  accompanied  by  an  effervescence,  or  rapid  evolu- 
tion of  gas  ;  and  second,  an  invisible  difference  ;  for  the  resulting 
liquid  is  not  a  solution  of  potassium  carbonate,  but  of  potassium 
nitrate.  In  the  first  case,  the  process  is  not  accompanied  by  any 
chemical  change  ;  the  operation  is  therefore  called  simple  solution : 
the  original  substance  is  present  in  the  liquid,  and  can  be  obtained 
in  its  former  state  by  evaporating  the  water.  In  the  second  case, 
the  process  is  distinguished  as  chemical  solution,  because  chemical 
action  took  place  between  the  substance  dissolved  and  the  solvent, 
and  the  original  substance  cannot  be  got  back  by  evaporating  the 
solvent. 

3.  Evaporation.— The  process  of  changing  from  the  liquid 
to  the  gaseous  or  vaporous  state  is  known  as  evaporation.  This 
operation  is  greatly  accelerated  by  the  application  of  heat.  When 
it  takes  place  without  the  aid  of  external  heat,  the  process  is  spoken 
of  as  spontaneous  evaporation. 

EXERCISE  3. — Pour  the  two  solutions  obtained  in  Exercise  2 
into  separate  porcelain  evaporating-dishes,  and  heat  them  gently 
by  means  of  a  Bunsen  with  a  "  rose  "  burner  (as  shown  in  Fig.  4). 
Continue  the  operation  until  all  the  liquid 
has  evaporated  away  and  a  dry  residue 
is  left.  This  is  called  evaporating  to  dry- 
ness.  As  the  condition  of  dryness  is  ap- 
proached, the  flame  must  be  turned  down 
more  and  more,  to  prevent  the  substance 
from  "sputtering."  Try  to  conduct  the 
operation  so  that  as  little  as  possible  of 
the  residue  is  lost  in  this  way. 

The  two  residues  may  now  be  exa- 
mined by  one  simple  test,  which  will 
prove  that  the  one  from  the  watery  solu- 
tion is  the  same  as  it  was  before  being 
dissolved,  and  that  the  other  is  quite  dif- 
ferent. Add  to  each  a  few  drops  of  dilute 
nitric  acid  :  the  first  dissolves  with  effer- 
vescence, as  did  the  original  potassium 
carbonate  ;  the  other  is  unacted  on  by  the  acid. 

Sometimes  it  is  necessary  to  carry  on  the  operation  of  evapora- 
tion more  carefully  than  can  be  done  by  heating  the  dish  in  the 
manner  described.  In  this  case  the  process  is  conducted  upon 
a  steam-bath.  Water  is  boiled  in  a  metal  vessel  (resembling  a 
saucepan),  and  the  evaporating-dish,  supported  by  a  metal  ring 
which  forms  the  cover,  is  heated  by  the  steam.  The  following 
exercise  is  a  case  in  point. 

EXERCISE  4.— Dissolve  some  crystals  of  ammonium  nitrate  in 
a  little  water ;  place  half  the  solution  in  a  dish,  and  evaporate  it 


FIG.  4. 


Preliminary  Exercises.  5 

over  a  rose  burner.     Evaporate  the  other  portion  in  a  dish  upon  a 
steam-bath.     Note  the  difference  in  the  results  in  the  two  cases. 

4.  Fusion  is  the  term  used  to  denote  the  process  of  changing 
a  substance  from  the  solid  to  the  liquid  state  by  the  action  of  heat. 
Thus,  when  lead  is  heated  it  enters  into  a  state  of  fusion,  or,  shortly, 
it  fuses  or  melts.  Fusion  must  not  be  confounded  with  solution. 
Chemical  action  often  takes  place  when  one  of  the  reacting  sub- 
stances is  in  a  condition  of  fusion,  which  is  incapable  of  taking 
place  when  they  are  only  in  solution.  For  example — 

EXERCISE  5. — Dissolve  a  small  piece  of  potassium  hydroxide 
(caustic potash]  in  water,  and  add  to  the  colourless  solution  a  minute 
quantity  of  powdered  manganese  dioxide.  No  chemical  action 
takes  place. 

Place  a  similar  piece  of  potassium  hydroxide  in  a  dry  test-tube, 
and  heat  it :  the  solid  fuses  to  a  colourless  liquid.  Drop  into  the 
fused  mass  a  few  particles  of  the  manganese  dioxide.  Chemical 
action  at  once  takes  place,  resulting  in  the  formation  of  the  deep 
green-coloured  compound,  potassium  manganate.  (This  reaction 
is  used  as  a  test  for  manganese  compounds.) 

EXERCISE  6. — Place  a  small  quantity  of  powdered  barium 
sulphate  in  a  test-tube,  add  water,  and  boil  for  a  minute  or  two. 
If  the  amount  of  barium  sulphate  is  quite  small,  it  will  be  easy  to 
see  that  practically  none  of  it  dissolves.  Allow  it  to  settle,  and  pour 
a  few  drops  of  the  liquid  upon  a  watch-glass,  and  set  it  to  evapo- 
rate to  dryness  on  a  steam-bath. 

Treat  another  similar  quantity  of  the  barium  sulphate  with 
dilute  hydrochloric  acid,  and  evaporate  a  few  drops  of  the  liquid 
in  the  same  way.  The  result  of  these  two  operations  will  prove 
that  barium  sulphate  is  insoluble  in  either  water  or  hydrochloric  acid. 

Next  dissolve  a  little  sodium  carbonate  in  water,  and  add  to  the 
clear  solution  a  few  particles  of  barium  sulphate  ;  boil  the  liquid, 
and  observe  that  no  change  takes  place. 

Now  carefully  mix  a  small  quantity  of  barium  sulphate  with 
about  five  times  as  much  sodium  carbonate  ;  place  the  powder  in 
a  platinum  crucible,  supported  on  a  pipe-clay  triangle  in  the  manner 
shown  in  Fig.  5,  and  heat  strongly  with  a  blowpipe.  When  the 
mass  has  been  in  complete  fusion  for  a  few  minutes,  allow  the 
crucible  to  cool.  Then  place  it  on  its  side  in  a  small  beaker  with 
a  little  water,  and  warm  gently.  The  mass  in  the  crucible  will  soon 
become  disintegrated,  some  of  it  dissolving,  while  a  part  remains 
undissolved.  Filter  the  liquid  as  in  Exercise  i,  washing  the  residue 
upon  the  funnel  until  the  filtrate  no  longer  restores  the  blue  colour 
to  reddened  litmus  paper.  Now  pour  a  few  drops  of  dilute  hydro- 
chloric acid  upon  the  residue  on  the  filter,  receiving  the  liquid 
which  passes  through  in  a  fresh  beaker  or  test-tube.  Observe  that 
effervescence  at  once  takes  place.  But  this  residue  cannot  be 
sodium  carbonate,  because  that  salt,  being  soluble  in  water,  has 


6 


Qualitative  Analysis. 


been  all  removed  ;  neither  can  it  be  barium  sulphate,  for  that  com- 
pound has  been  shown  to  be  insoluble  in  dilute  hydrochloric  acid. 
By  the  process  of  fusion,  the  sodium  carbonate  and  barium  sulphate 
underwent  a  chemical  reaction,  resulting  in  the  formation  of  sodium 
sulphate  and  barium  carbonate.  The  former  salt,  being  soluble  in 
water,  was  dissolved  in  that  liquid  along  with  the  excess  of  sodium 
carbonate.  The  barium  carbonate  is  insoluble  in  water,  but  dis- 
solves in  dilute  hydrochloric  acid,  forming  barium  chloride  (soluble) 

and  carbon  dioxide,  which 
escapes  as  gas.  Therefore, 
by  fusion  the  insoluble 
barium  siilphate  is  converted 
into  soluble  barium  carbo- 
nate. 

5.  Precipitation. 
— When  chemical  action 
takes  place  between  sub- 
stances in  solution,  and  one 
of  the  products  of  the  action 
is  insoluble,  the  latter  sub- 
stance is  thrown  out  of  solu- 
tion, or  precipitated.  The 
substance  so  thrown  down  is 
termed  a  precipitate. 

EXERCISE  7. — Dissolve 
a  minute  quantity  of  sodium 
chloride  (common  salt)  in 
water  in  a  test-tube.  In 
another  tube  dissolve  a  small 
crystal  of  silver  nitrate,  and 
mix  the  two  solutions  to- 
gether. The  two  compounds 
react  upon  each  other,  form- 
ing sodium  nitrate  (soluble 
in  water)  and  silver  chloride 
The  insoluble  white  precipitate  is  therefore 


FIG.  5. 

(insoluble  in  water), 
the  silver  chloride. 


If,  in  the  above  example,  the  two  substances  are  mixed  in  a 
particular  proportion,  there  may  have  been  exactly  the  amount  of 
sodium  chloride  necessary  to  supply  chlorine  enough  to  unite  with 
the  whole  of  the  silver  in  the  silver  nitrate  used.  In  this  case  there 
would  be  nothing  left  in  sohttion  but  sodium  nitrate,  i.e.  no  excess 
of  either  silver  nitrate  or  of  sodium  chloride.  Ascertain  if  this 
happened  to  be  the  case  in  Exercise  7,  by  the  following  experi- 
ment : — 


Preliminary  Exercises.  7 

EXERCISE  8.— Filter  the  mixture  obtained  above,  and  divide 
the  nitrate  into  two  portions.  To  one  add  a  single  drop  of  a 
solution  of  sodium  chloride,  (i)  If  a  precipitate  is  formed,  it 
proves  that  there  is  some  silver  nitrate  present,  and  that  therefore 
an  excess  of  this  compound  was  used  in  Exercise  7.  Continue 
adding  the  sodium  chloride  solution  one  drop  at  a  time,*  shaking 
or  stirring  the  liquid  after  every  addition,  so  long  as  it  produces 
further  precipitation. 

(2)  If  no  precipitate  is  thrown  down  by  adding  sodium  chloride, 
add  to  the  second  portion  of  the  filtrate  a  single  drop  of  silver  nitrate 
solution.  If  this  gives  a  precipitate,  it  proves  that  sodium  chloride 
is  present,  and  that  therefore  an  excess  of  this  substance  had  been 
employed  in  Exercise  7.  Continue  adding  the  silver  solution  drop 
by  drop,  with  constant  stirring,  so  as  to  hit  off  as  nearly  as  possible 
the  exact  point  when  it  just  ceases  to  produce  any  further  pre- 
cipitate. 

The  exact  point  at  which  precipitation  is  complete  is  not  equally 
easy  to  determine  in  all  cases.  Some  precipitates  are  heavy, 
granular,  or  crystalline,  and  settle  quickly  ;  others  again  are  light 
or  flocculent,  and  only  subside  slowly  and  imperfectly,  so  that  it  is 
difficult  to  see  whether  the  addition  of  more  of  one  of  the  solutions 
does  or  does  not  produce  any  additional  precipitate.  In  such  cases 
the  liquid  should  be  filtered,  and  the  filtrate  tested  by  adding  a  few 
drops  more  of  the  precipitant. 

Very  often  several  substances  present  together  in  one  and  the 
same  liquid,  form  insoluble  compounds  with  another  which  is 
added.  These  will  not  be  all  precipitated  simultaneously,  but  in  a 
certain  order  one  after  the  other,  the  precipitation  of  one  being 
more  or  less  complete  before  that  of  the  next  begins.  The  substance 
being  added,  first  selects  the  compound  present  for  which  it  has 
the  greatest  chemical  affinity,  and  afterwards  that  with  which  it 
unites  less  eagerly.  This  being  so,  it  will  be  evident  that  unless 
care  be  taken  to  ensure  complete  precipitation,  it  might  easily 
happen  that  the  whole  of  one  of  the  substances  present  in  the 
solution  escaped  precipitation.  It  is  of  the  utmost  importance, 
therefore,  in  analysis,  to  be  quite  sure  that  precipitation  is  as  com- 
plete as  possible.  On  the  other  hand,  the  reckless  addition  of  pre- 
cipitants  is  a  fault  which  must  be  as  carefully  guarded  against,  as 
it  is  almost  as  fruitful  a  source  of  trouble  as  the  other. 

In  most  instances,  also,  it  is  essential  to  wash  the  precipitate 
until  it  is  quite  free  from  any  of  the  soluble  substances  present  in 

*  When  solutions  are  to  be  added  drop  by  drop,  it  is  best  to  use  a  pipette ; 
that  is,  a  piece  of  ordinary  glass  tube  drawn  to  a  point  at  one  end,  and  about 
6  or  8  inches  long. 


8  Qualitative  Analysis. 

the  liquid,  as  explained  in  Exercise  i.  A  precipitate  may  be  re- 
moved from  the  filter  either  by  means  of  a  spatula  (preferably 
platinum,  but,  failing  this,  either  glass  or  porcelain ;  iron  should 
never  be  used),  or  by  pushing  a  glass  rod  through  the  apex  of  the 
filter,  and  then  washing  the  precipitate  through  by  means  of  the 
wash-bottle,  or  by  dissolving  it  off  by  pouring  into  the  funnel  the 
liquid  to  be  used  for  its  solution. 

EXERCISE  9. — Add  a  solution  of  sodium  carbonate  to  a  solution 
of  barium  chloride,  until  precipitatioij  is  just  complete.  Barium 
carbonate  is  precipitated,  and  sodium  chloride  remains  in  solution. 
Pour  the  mixture  upon  three  separate  filters,  and  wash  the  pre- 
cipitate on  each  until  quite  free  from  sodium  chloride  (see  Exercises 
7  and  8),  getting  the  precipitate  well  down  into  the  apex. 

Take  the  first  filter,  and  remove  a  portion  of  the  precipitate  with 
a  spatula.  If  the  quantity  in  the  funnel  is  small,  then  carefully  draw 
the  paper  cone  out  of  the  funnel,  spread  it  open  upon  a  flat  sheet  of 
glass,  and  scrape  off  as  much  of  the  precipitate  as  possible  with  the 
spatula,  and  transfer  it  to  a  test-tube.  Dissolve  it  by  adding  a  few 
drops  of  hydrochloric  acid. 

Through  the  apex  of  the  second  filter  push  a  glass  rod,  and  wash 
the  precipitate  through  into  a  test-tube,  using  a  fine  jet  of  water, 
and  as  little  of  it  as  possible.  Dissolve  this  also  by  adding  a  few 
drops  of  the  same  acid. 

Upon  the  third  filter  pour  a  small  quantity  of  hydrochloric  acid, 
collecting  the  filtrate  in  a  test-tube.  Pour  the  filtrate  back  over 
the  precipitate  once  or  twice,  until  the  whole  has  dissolved. 

6.  Ignition. — Strictly  speaking,  this  word  carries  with  it  the 
idea  of  combustion.  In  common  speech  it  signifies  the  act  of 
"  setting  fire  "  to  an  inflammable  substance  ;  and  in  more  scientific 
language  we  speak  of  the  ignition  temperature,  or  the  igniting- 
point  of  a  body,  meaning  thereby  the  temperature,  to  which  it  is 
necessary  to  raise  it  in  order  that  combustion  may  be  initiated. 
Unfortunately,  in  analytical  phraseology  the  term  ignition  is  used 
in  a  somewhat  slipshod  way  to  denote  a  variety  of  operations 
where  substances  are  simply  strongly  heated,  and  where  the  idea 
of  combustion  is  altogether  excluded.  In  this  book  the  words  heat 
or  strongly  heat  will  be  used  instead  of  ignite  to  signify  these 
operations. 

EXERCISE  10.  Strongly  heating  in  an  open  dish— -Place  a 
little  solution  of  ammonium  chloride  in  a  small  evaporating-dish, 
and  evaporate  to  dryness.  Then  strongly  heat  the  dish  with  the 
dry  residue  until  no  more  white  fumes  (consisting  of  the  volatilising 
ammonium  chloride)  are  evolved.  If  the  dish  has  been  heated  all 
over  there  should  then  be  nothing  left  in  it.  The  complete 
vaporisation  of  the  salt  is  more  quickly  and  certainly  accomplished 


Preliminary  Exercises.  9 

by  using  a  small  platinum  capsule  or  crucible  in  which  to  heat  the 
residue  obtained  by  evaporating  the  solution  to  dryness. 

EXERCISE  n.  Strongly  heating  in  a  tube  closed  at  one  end. — 
Place  a  minute  quantity  of  mercuric  oxide  in  a  small  test-tube 
(4  inches  x  j^),  and  apply ,heat  to  the  compound.  Note  the  change 
of  colour.;  also  that  it  gradually  disappears,  and  that  a  sublimate 
collects  on  the  cool  part  of  the  tube,  having  a  white  metallic 
appearance.  Test  the  evolved  oxygen  by  means  of  a  glowingisplint 
of  wood.  By  means  of  a  paper  "  spill "  rub  the  metallic  sublimate, 
and  (if  necessary,  with  a  pocket  lens)  see  the  globules  of  mercury. 

EXERCISE  12.  Heating  in  the  blowpipe  flame. — Select  a  piece 
of  small  tubing  of  lead  glass,  and  heat  it  in  a  blowpipe  flame, 
holding  the  glass  in  the  extreme  tip  of  the  flame  until  it  is  red  hot. 
Then  gradually  bring  it  further  into  the  flame,  and  observe  that 
when  the  glass  reaches  the  inner  cone  of  the  flame  a  film  begins  to 
appear  upon  the  red-hot  portion.  On  withdrawing  the  glass  to  the 
tip  of  the  flame  again,  this  film  gradually  disappears.  Bring  the 
glass  once  more  into  the  inner  cone  of  flame,  and  when  the  film 
has  again  made  its  appearance,  remove  the  glass  and  allow  it  to 
cool.  It  will  then  be  seen  that  what  appeared  like  a  film  when  it 
was  hot,  is  a  black  shining  metallic-looking  deposit  in  the  glass. 
This  deposit  is  metallic  lead.  The  lead  compound  in  the  glass, 
when  heated  in  the  inner  cone  of  flame,  is  reduced  to  the  metallic 
state  ;  and  when,  after  being  so  reduced,  it  is  heated  in  the  tip  of 
the  flame  (i.e.  in  the  outer  cone  or  sheath  of  the  flame),  the  metal  is 
again  oxidised.  The  inner  flame  is  therefore  called  the  reducing 
flame,  and  the  outer  cone  is  distinguished  as  the  oxidising  flame.* 

EXERCISE  13.  Heating  on  charcoal  in  the  blowpipe  flame. — 
Select  a  close-grained  piece  of  charcoal,  as  free  as  possible  from 
cracks,  and  file  a  flat  surface  upon  it  with  a  broad,  flat  file.t  On 
the  flat  part  scoop  a  small  hollow,  and  place  in  it  a  little  red-lead 
mixed  with  about  an  equal  quantity  of  sodium  carbonate.  Heat 
this  mixture  in  the  inner  blowpipe  flame,  holding  the  blowpipe 
and  the  charcoal  in  the  manner  shown  in  Fig.  6,  so  that  the  flame 
shall  play  along  the  surface  of  the  charcoal.  Very  quickly  the  lead 
oxide  will  be  reduced  to  the  metallic  state,  and  appear  in  the  form 
of  brilliant  silvery  globules.  When  the  charcoal  is  removed,  it 
will  be  seen  that  surrounding  the  cavity  there  is  a  yellowish  deposit, 
or  incrustation.  This  consists  of  lead  oxide.  If  the  outer  tip  of 
the  flame  be  directed  upon  this  incrustation  it  will  quickly  disappear 
and  will  impart  a  bluish  colour  to  the  end  of  the  flame. 

Pick  out  one  or  two  of  the  globules  of  metal,  and  gently  strike 
one  with  a  small  hammer,  or  with  a  pestle,  upon  some  hard  surface. 

*  The  memory,  of  the  beginner  may  be  aided  by  the  alliteration,  Outer 
Oxidising.  The  inner  flame  is  a  reducing  agent  by  reason  of  the  fact  that 
within  the  cone  there  is  an  excess  of  strongly  heated  coal-gas  ;  whereas  in  the 
outer  flame  there  is  an  excess  of  heated  atmospheric  oxygen. 

f  Specially  prepared  rectangular  blocks  of  charcoal  (6  inches  long  and  i 
square  inch  section)  are  sold  for  the  purpose.  One  such  block  can  be  used 
many  times. 


10 


Qualitative  A  nalysis. 


Note  whether  the  metal  is  hard  and  brittle,  or  soft  and  malleable. 
Also  further  identify  the  metal  as  lead  by  rubbing  it  upon  a  piece 
of  paper,  which  will  be  marked  by  it  much  as  by  an  ordinary 
pencil. 

EXERCISE  14. — Heat  on  another  piece  of  charcoal  a  crystal  or 
two  of  zinc  sulphate  with  a  little  sodium  carbonate  in  the  inner 
blowpipe  flame.  No  metallic  globules  are  formed  in  this  case, 
because  zinc  is  too  easily  oxidised ;  but  an  incrustation  appears  on 
the  charcoal,  which  is  canary-yellow  while  hot,  but  turns  white  on 
cooling.  Touch  the  incrustation  with  a  single  drop  of  a  solution 
of  cobalt  nitrate,  and  again  heat  it,  using  the  outer  flame.  The 


FIG.  6. 

incrustation  then  becomes  green.  Notice  that  the  incrustation  is 
not  driven  off  by  being  thus  heated,  because  zinc  oxide  is  not 
volatile. 

7.  Fusion,  with  Borax. — When  borax  is  strongly  heated,  it 
melts  to  a  clear  vitreous  mass.  In  this  condition  it  is  capable  at  a 
high  temperature  of  dissolving  many  metallic  compounds,  giving  in 
some  cases  characteristically  coloured  glasses. 


Preliminary  Exercises*  n 

EXERCISE  15. — Twist  the  end  of  a  piece  of  platinum  wire  into 
a  small  round  loop  or  eye,*  and  pick  up  a  little  borax  upon  it  by  first 
heating  the  wire  and  then  dipping  it  while  hot  into  the  powdered 
salt.  On  heating  the  borax  upon  the  wire  in  a  blowpipe  flame,  it 
first  swells  up,  and  finally  fuses,  forming  a  transparent  colourless 
bead  of  borax  glass.  Allow  the  bead  to  cool,  and  touch  it  with  a 
glass  rod  which  has  been  dipped  into  a  solution  of  any  cobalt  salt, 
so  as  to  bring  only  a  minute  quantity  of  the  cobalt  compound  upon 
the  bead.  Heat  the  bead  once  more,  and  notice  that  as  it  melts 
the  borax  loses  its  transparent  appearance.  When  again  allowed 
to  cool,  the  bead  will  appear  of  an  azure  blue  colour. 

If  too  much  of  the  cobalt  salt  was  employed,  the  bead  may 
appear  almost  black  ;  in  this  case  a  part  of  it  may  be  shaken  off 
when  it  is  fluid,  and  more  borax  picked  up  and  melted  with  what 
remains  of  the  original  bead  upon  the  wire.  If  too  little  cobalt  is 
present,  the  colour  will  be  correspondingly  pale.  The  colour  of 
the  bead  is  best  examined  by  holding  it  against  a  white  object 
(such  as  the  bottle  of  borax  itself)  in  a  good  light. 

Fuse  the  bead  again,  holding  it  first  in  the  outer  flame,  and 
afterwards  in  the  inner  flame,  and  see  that  in  each  case  when 
cold  the  blue  colour  remains  the  same. 

EXERCISE  16. — Make  another  borax  bead,  and  touch  it  with 
a  small  quantity  of  a  solution  of  manganous  sulphate.  Heat  this 
in  the  outer  blowpipe  flame.  After  cooling,  examine  the  colour 
carefully.  Pale  violet,  lilac,  purple,  or  amethyst.  Heat  the  bead 
again,  holding  it  in  the  inner  flame.  Notice  that  it  gradually 
loses  its  opacity  ;  that  as  it  is  heated,  something  in  the  fused  mass 
which  seems  to  give  it  an  appearance  of  muddiness  clears  away, 
and  the  molten  globule  looks  clear.  When  it  is  in  this  condition 
remove  it,  and  when  cold  it  will  be  found  to  have  lost  its  colour 
entirely.  Manganese  compounds  there- 
fore give  a  purplish  bead  in  the  outer 
flame,  which  becomes  colourless  upon 
being  heated  in  the  reducing  flame. 

8.  Neutralisation. — When  an  acid 
is  carefully  mixed  with  an  alkali  (the 
substances  being  in  solution),  a  point  is 

*  For  greater  convenience,  as  well  as  economy, 
a  short  piece  of  wire  (about  2  inches)  should  be 
fixed  into  a  glass  tube,  about  the  same  length, 
to  serve  as  a  handle.  The  glass  tube  is  first  drawn 
out  to  a  point,  and  the  wire  inserted  into  the  fine 
end.  On  bringing  this  into  a  blowpipe  flame, 
the  glass  fuses  round  the  wire  and  holds  it. 
Two  or  three  of  these  should  be  made,  and  a 
convenient  plan  is  to  fit  the  glass  tube  into  a  VIG.  7« 

cork,  so  that  when  not  actually  in  use  the  wires 

can  be  kept  in  small  test-tubes  containing  dilute   hydrochloric   acid    as  in 
Fig.  7. 


12  Qualitative  Analysis. 

reached  when  the  mixture  no  longer  possesses  the  properties  of 
either  the  acid  or  the  alkali.  The  solution  is  then  said  to  be 
neutral.  The  point  of  neutrality  is  ascertained  by  the  use  of 
certain  sensitive  colouring  matters  which  have  their  colour  changed 
by  acids  and  alkalies.  The  commonest  of  these  is  litmus,  the 
solution  of  which  in  water  has  a  purple  colour,  capable  of  being 
turned  red  by  acids,  and  blue  by  alkalies.  The  yellow  colour  of 
turmeric  is  changed  to  brown  by  alkalies,  but  is  not  altered  by 
acids,  therefore  this  can  only  be  used  to  indicate  alkalinity,  and 
will  not  discriminate  between  a  neutral  and  an  acid  liquid. 

EXERCISE  17. — Add  a  few  drops  of  litmus  solution  to  a  little 
dilute  hydrochloric  acid  in  a  beaker  standing  upon  a  piece  of  white 
paper,  or  a  white  tile.  Add  to  the  red  liquid  some  solution  of 
sodium  hydroxide,  adding  it  cautiously  in  small  quantities,  with 
constant  stirring,  until  the  colour  of  the  litmus  is  just  turned  blue. 
The  liquid  is  now  alkaline.  By  means  of  a  glass  rod  moistened 
with  the  dilute  acid,  introduce  a  minute  additional  quantity  of  the 
acid,  so  as  to  cause  the  colour  of  the  litmus  to  become  of  a  purple 
tint.  The  solution  is  then  neutral,  and  the  least  trace  of  either 
acid  or  alkali  will  at  once  turn  it  red  or  blue,  as  the  case  may  be. 
(Instead  of  adding  litmus  solution,  papers  tinted  with  litmus  may 
be  dipped  into  the  liquid.) 

All  substances  which  redden  litmus  are  not  acids,  although  all 
acids  will  redden  litmus.  That  is  to  say,  there  are  many  things 
which  have  an  acid  or  sour  taste  which  do  not  belong  to  that  class 
of  compounds  which  chemists  call  adds.  An  acid  may  be  denned 
as  a  compound  containing  hydrogen  which  can  be  displaced  by  a 
metal,  when  the  latter  is  presented  to  it  in  combination  as  a 
hydroxide.  The  hydrogen  which  is  thus  displaced  is  not  liberated 
as  free  hydrogen,  but  unites  with  the  hydrogen  and  oxygen  of  the 
metallic  hydroxide  to  form  water,  while  the  metal  takes  the  place 
of  the  hydrogen  thus  displaced  from  the  acid. 


CHAPTER  II. 
ANALYTICAL  CLASSIFICATION. 

THE  word  analysis,  in  its  strict  meaning,  signifies  the  breaking 
up  or  separation  of  a  compound  substance  into  its  constituent 
parts.  It  is  the  true  antithesis  of  the  word  synthesis,  which 
means  the  building  up  of  a  compound  from  its  constituents. 

But  the  word  analysis  has  come  to  bear  a  wider  meaning, 
and  to  include  all  the  various  processes  and  operations  which 
chemists  make  use  of  in  order  to  find  out  what  any  compound  is 
composed  of,  or  to  enable  them  to  identify  the  substance,  quite 
irrespective  of  whether  or  not  the  process  involves  the  breaking  up 
of  the  body  into  its  component  parts.  Thus,  a  chemist  will  often 
recognise  a  substance  by  its  particular  crystalline  form,  or  from 
some  other  characteristic  appearance  it  may  present  when  examined 
under  a  microscope  (microscopic  analysis).  Or  sometimes  he  can 
detect  the  presence  of  certain  elements  in  an  unknown  substance, 
by  examining  the  light  which  is  emitted  when  the  compound  is 
strongly  heated  (spectrum  analysis). 

Reactions. — Most  analytical  operations,  however,  involve 
some  chemical  change.  These  changes  are  called  reactions. 
When  the  change  is  effected  by  strongly  heating  the  substance, 
it  is  described  as  a  dry  reaction,  or  a  reaction  in  the  dry  way. 
This  is  to  distinguish  this  class  of  reactions  from  those  which  take 
place  between  substances  that  are  dissolved,  either  in  water  or 
some  other  liquid,  and  which  are  sometimes  spoken  of  as  ivet 
reactions,  or  reactions  in  the  wet  way. 

Most  analytical  reactions  are  "  double  decompositions,"  in  which 
one  of  the  products  of  the  chemical  action  is  either  markedly 
different  from  the  others  and  from  the  reacting  compounds,  in  its 
solubility,  or  its  colour  ;  or  where  it  is  evolved  as  a  gas  having 
properties  by  which  it  may  be  readily  identified.  For  instance,  the 
two  compounds  barium  chloride  and  sodium  sulphate  are  soluble 
in  water,  forming  colourless  solutions  ;  if  these  are  mixed  together, 
"  double  decomposition  "  takes  place,  resulting  in  the  formation  of 
sodium  chloride  and  barium  sulphate,  thus — 

Bad,  +  Na2SO4  =  2NaCl  +  BaSO, 


14  Qualitative  Analysis. 

The  barium  sulphate  is  practically  insoluble  in  water,  and  con- 
sequently is  precipitated.  Now,  if  we  know  some  property  belong- 
ing to  this  precipitate  of  barium  sulphate  which  is  so  characteristic 
of  the  compound  that  we  could  thereby  identify  it  and  distinguish 
it  from  all  other  white  precipitates,  then  this  reaction  between 
barium  chloride  and  sodium  sulphate  can  obviously  be  used  as  a 
means  of  testing  for  the  presence  of  either  a  soluble  barium  salt 
or  a  soluble  sulphate.  For  if,  on  adding  a  solution  of  sodium 
sulphate  to  an  unknown  solution,  barium  sulphate  were  precipitated, 
the  unknown  liquid  must  have  contained  a  soluble  barium  salt ;  or, 
on  the  other  hand,  if  we  add  barium  chloride  to  an  unknown  solu- 
tion and  obtain  barium  sulphate  again,  then  this  unknown  solution 
must  have  contained  a  sulphate.* 

Reagents. — The  materials  that  are  used  to  bring  about 
analytical  reactions  are  termed  reagents.  Thus,  in  the  illustration 
given  above,  the  sodium  sulphate  is  the  reagent  when  it  is  added 
to  the  unknown  solution  in  order  to  test  for  barium  ;  while  the 
barium  chloride  is  the  reagent  when  it  is  used  to  test  for  a  sulphate. 
Some  reagents  are  capable  of  causing  reactions  of  a  similar 
character  with  a  number  of  substances  ;  such  are  often  known  as 
general  reagents.  Others,  again,  are  employed  because  they  pro- 
duce a  characteristic  reaction  with  some  one  substance  in 
particular  ;  these  are  distinguished  as  special  reagents. 

Reagents  are  the  tools  with  which  the  analyst  works,  and  upon 
the  intelligent  and  skilful  use  of  them  everything  depends.  In  most 
laboratories  the  student  finds  himself  supplied  with  all  the  necessary 
reagents  ready  prepared  ;  but  for  the  help  of  those  who  may  require 
to  make  them  up,  brief  directions  for  doing  so  are  given  in  the 
Appendix. 

Analytical  Classification.— Substances  are  usually  divided 
into  two  classes,  namely,  (i)  Metals,  and  (2)  Acid-radicals. 
These  are  also  sometimes  called  positive  radicals  and  negative 
radicals  respectively.  When  analysing  such  a  compound  as  sodium 
chloride,  NaCl,  the  sodium  and  the  chlorine  are  each  separately 
detected  :  the  sodium  is  the  metal  (or  positive  radical),  and  the 
chlorine  is  the  acid  (or  negative)  radical.  But  in  such  a  case  as 
sodium  nitrate,  NaNO3,  we  do  not  separately  detect  the  sodium, 
the  nitrogen,  and  the  oxygen,  but  the  sodium  and  the  negative  or 
acid  radical  represented  by  the  formula  NO3.  Or,  again,  when 
such  a  compound  as  ammonium  sulphate,  (NH4)2SO4,  is  submitted 
to  analysis,  we  do  not  separately  test  for  the  elements  nitrogen, 

*  Sulphuric  acid  being  included,  as  hydrogen  sulphate. 


Analytical  Classification.  15 

hydrogen,  sulphur,  and  oxygen,  but  for  the  positive  radical  NH4, 
and  the  acid-radical  SO4. 

Sometimes  the  radicals,  whether  metals  or  acid  radicals,  may  be 
detected  by  being  actually  isolated,  in  which  case  they  are  recog- 
nised by  their  known  properties  in  the  free  state.  For  example, 
from  the  compound  lead  chloride,  PbCl2,  it  is  easy  to  isolate  both 
the  lead  and  the  chlorine.  The  metal  lead  so  obtained  is  readily 
identified  by  its  familiar  physical  properties,  while  the  gas  chlorine  is 
equally  easily  distinguished  by  its  own  well-known  characteristics. 

In  some  cases,  where  a  radical  is  incapable  of  isolated  existence, 
it  may  be  detected  by  the  separation  of  some  product  of  its  decom- 
position. Thus,  in  such  a  compound  as  ammonium  carbonate, 
(NH4)2CO3,  neither  the  positive  radical  NH4,  nor  the  acid-radical 
CO3  can  exist  in  the  free  or  uncombined  state.  But  we  detect  the 
presence  of  the  former  by  the  evolution  of  ammonia,  NH3,  and  the 
latter  by  the  expulsion  of  carbon  dioxide,  CO2,  from  the  compound. 

In  the  large  majority  of  cases,  however,  whether  the  various 
radicals  are  capable  of  isolated  existence  or  not,  they  are  detected 
by  causing  them  to  pass  into  fresh  combinations  with  certain 
reagents,  whereby  new  compounds  are  formed  which  are  readily 
recognised  by  their  known  properties.  Thus,  in  the  case  of  sodium 
chloride  above  quoted,  instead  of  isolating  the  chlorine,  we  can 
employ  the  reagent  silver  nitrate,  AgNO3.  When  this  is  added  to 
a  solution  of  sodium  chloride,  double  decomposition  takes  place, 
and  silver  chloride,  AgCl,  is  formed,  which,  being  insoluble  in  water, 
is  precipitated.  Silver  chloride  has  properties  by  which  it  is  easily 
recognised,  hence  by  the  formation  of  this  compound  we  can  detect 
the  presence  of  the  chlorine  in  sodium  chloride. 

In  all   such   cases   as   these   the   interaction   is  between   the 

ions  into  which  the   compounds    dissociate   when   dissolved    in 

water.      A  solution   of   sodium   chloride,    for   example,  contains 

+  + 

Na  and  Cl  ions  ;  the  silver  nitrate  contains  Ag  and  NO3  ions. 

When  these  solutions  are  mixed,  the  positive  silver  ions  unite  with 
the  negative  chlorine  ions  to  produce  the  electrically  neutral  and 
insoluble  silver  chloride,  which  therefore  separates  out.  The  silver 
ions,  therefore,  are  the  test  for  chlorine  ions,  and  vice  versa 
chlorine  ions  are  the  reagents  for  detecting  silver  ions.  Any 
chlorine  compound  which  on  dissociation  furnishes  chlorine  ions, 
will  therefore  respond  to  this  test  with  silver  ions.  There  are, 
however,  many  compounds  containing  chlorine  which  give  no 
precipitate  of  silver  chloride  on  the  addition  of  silver  nitrate. 
Familiar  among  these  are  the  chlorates  and  perchlorates.  These 


16 


Qualitative  A  nalysis. 


compounds  dissociate  on  solution,  not  into  simple  chlorine  ions, 
but  into  the  complex  C1O3  and  C1O4  ions.  Such  solutions,  there- 
fore, contain  no  Cl  ions  and  are  therefore  incapable  of  forming  AgCl 

+ 
with  Ag  ions. 

In  analytical  classification  the  term  "  metal  "  includes,  besides 
the  metals  proper,  certain  metalloidal  elements,  such  as  arsenic, 
selenium,  and  others  (which,  strictly  speaking,  are  not  true  metals, 
but  which  lie  on  the  borderland  between  the  metals  and  the  non- 
metals),  and  also  the  compound  positive  radical  ammonium,  NH4. 
These  "  metals,"  then,  are  divided  into  a  number  of  groups,  based 
on  the  behaviour  of  their  compounds  towards  certain  reagents. 

As,  however,  more  than  one  scheme  or  plan  of  analysis  is 
possible,  due  either  to  a  different  choice  of  reagents  or  to  their  use 
in  a  different  order,  so  the  analytical  classification  of  the  metals 
which  is  followed  by  some  chemists  varies  somewhat  from  that 
used  by  others. 

In  this  book  the  following  arrangement  will  be  adopted  : — 


Group  I. 

Group  II. 

Group  III. 

Group  IV. 

Group  V. 

Silver 

Mercury 

Aluminium 

Barium 

Ammonium 

Mercury 

Lead                  Chromium 

Strontium 

Sodium 

Lead* 

Bismuth           Iron 

Calcium 

Potassium 



Copper             Nickel 



Magnesium 

Thallium  4* 

f!  *"i  rl  in  1  1  1  TT1              f*AVk*»14* 

Tungsten 

V/&1UII11  UU1               OODillL 

Antimony       Manganese 

Lithium 

Arsenic            Zinc 

Rubidium 

Tin 



Caesium 

Gold 

Beryllium 

Platinum 

Zirconium 



Thorium 

Ruthenium 

Cerium 

Rhodium 

Scandium 

Palladium 

Yttrium 

Osmium 

Lanthanum 

Iridium 

Ytterbium 

Tellurium 

Titanium 

Selenium 

Tantalum 

Molybdenum 

Niobium 
Uranium 

Indium 

Thallium 

Vanadium 

*  The  reason  why  certain  metals  are  placed  in  more  than  one  division  will 
appear  later. 

f  The  substances  printed  in  small  type  are  usually  called  rare  elements. 


Analytical  Classification.  17 

In  the  regular  course  of  analysis,  the  groups  are  separated  in 
the  order  in  which  they  are  here  numbered.  The  "  silver,  lead, 
mercury  "  group  is  separated  first,  and  the  "ammonium,  sodium," 
etc.,  group  last.  In  studying  the  reactions  of  the  metals,  however, 
it  is  more  usual  to  begin  with  the  fifth  group,  for  the  reason  that 
the  compounds  of  these  metals  are  less  complex,  and  the  student 
is  therefore  led  on  gradually  from  what  is  comparatively  simple 
to  that  which  is  more  difficult.  Because  of  this,  some  chemists 
prefer  to  number  the  groups  in  the  reverse  order,  that  is,  in  the 
order  in  which  the  preliminary  sttidy  of  them  is  made,  instead  of 
the  order  in  which  they  are  actually  disposed  of  in  the  course  of 
analysis. 

The  reagents  by  means  of  which  the  elements  are  separated 
into  these  groups,  and  which  are  known  as  group  reagents,  must 
be  used  in  regular  order.  Each  is  only  capable  of  separating  its 
own  family  of  metals  from  those  coming  after  it  in  the  series,  and 
not  those  going  before.  For  example,  the  group-reagent  for 
Group  II.  is  only  capable  of  separating  the  metals  of  this  family 
from  those  of  III.,  IV.,  and  V.,  but  not  from  those  of  Group  I. 
If,  therefore,  the  metals  of  Group  I.  are  not  first  separated,  they  may 
be  precipitated  along  with  the  members  of  the  second  family  by 
the  group-reagent  for  that  family. 

The  various  group-reagents  (or  general  reagents)  and  the 
particular  compounds  of  the  metals  which  are  precipitated  by  them 
are  indicated  in  the  following  table  : — 

GROUP     I.  —  General    reagent,    Hydrochloric     acid,    HC1, 

precipitates 

AgCl,  Hg!2Cl2f  PbCl2  (partially  soluble). 
1,  H2WO4.)* 


GROUP  \\.-General  reagent,  Sulphuretted  hydrogen,  H2S, 

precipitates  in  acid  solution 

"PbS,  Hg  S,  Bi  S  ,  CuS,  CdS         \  Insoluble  in  ammo- 
(Ru2S3,  Rh2S3,  PdS,  OsS.)      /     nium  sulphide. 


OOf 


Sb2S3,  As2S3,  Sn"S  and  Sn*S2,  AuS,  PtS 


Soluble  in 

ammon. 
(Ir2S3,  TeS2(  be,  Mob3.)  sulphide. 

*  The  compounds  represented  by  the  formulae  in  small  type  are  those  of 
the  so-called  rare  elements.  They  are  included  in  the  table  in  order  to  give,  in 
a  bird's-eye  view  as  it  were,  not  only  their  position  in  the  scheme  of  classifica- 
tion, but  also  the  composition  of  their  compounds,  which  are  precipitated  by 
the  group-reagents. 

C 


1 8  Qualitative  Analysis. 

GROUP  III. — General  reagent,  Ammonium  sulphide,  (NHJJS, 
precipitates  in  presence  of  ammonium  chloride  and  ammonia 
(«)  Hydrated  I  A12(HO),,  Cr2(HO),. 


compounds       (Be(HO)2l  Zr(HO)4,  Th(HO)4,  Ce(HO)3.  Sc(HO)3,  Y(HO)3, 
La(HO)3,  Yb(HO)3,  H2TiO3>  H3TaO4(  H3NbO4.) 

|  FeS,  NiS,  CoS,  MnS,  ZnS. 

(b}  Sulphides  j  (UO2S,  InS,  T12S)  (Vanadium 

converted  into  soluble  ammonium  thiovanadate  *). 

GROUP  IV.  —  General  reagent.  Ammonium  carbonate, 
(NHJ.CO  ;,  precipitates  in  presence  of  ammonium  chloride 
and  ammonia 

BaC03,  SrCO,,  CaCO3. 

GROUP     V.  —  No   general   reagent.      The    group    consists    of 
(NH4),  Na,  X,  'Mg.  (Li,  Rb,  Cs.) 


*  Vanadium  belongs  to  the  "arsenic  and  antimony"  family  in  the  natural 
classification  of  the  elements,,  The  sulphide  is,  however,  not  precipitated  by 
H2S,  V2O5  being  thereby  reduced  to  V2O4,  which  gives  a  blue  colour  to  the 
liquid.  NH4C1,  in  presence  of  ammonia,  precipitates  white  ammonium  meta- 
vanadate,  NH4VO3,  but  ammonium  sulphide  converts  this  into  the  soluble 
ammonium  thiovanadate,  which  gives  a  brown  colour  to  the  solution.  The 
true  group-reagent,  therefore,  does  not  actually  precipitate  this  metal. 


CHAPTER   III. 
REACTIONS  OF  THE  METALS  OF  GROUP  V. 

THIS  group  contains  the  alkali  metals  (ammonium  being  regarded 
as  a  metal),  and  also  the  element  magnesium,  which  is  more  nearly 
allied  to  the  metals  of  the  alkaline  earths.  The  members  of  this 
family  are  not  precipitated  by  any  group-reagent,  but  they  are 
(with  the  exception  of  ammonium)  separately  tested  for  in  the 
solution  which  is  obtained  after  the  metals  of  Groups  I.  to  IV.  have 
been  removed.  By  referring  to  the  table  on  p.  17,  it  will  be  seen 
that,  in  the  course  of  separating  the  various  groups,  certain  am- 
monium compounds  are  employed,  therefore  it  will  be  obvious  that 
it  is  necessary  to  test  for  this  "  metal  "  in  the  substance  under  ex- 
amination before  adding  any  ammoniacal  compounds. 


Ammonium, 

DRY  REACTIONS.  —  When  heated  alone  in  a  glass  tube,  ammo- 
nium salts  undergo  change. 

(a)  If  the  acid  is  readily  volatile,  the  salt  dissociates,  but  the 
ammonia  and  the  volatile  acid,  as  they  together  pass  away  from 
the  heated  area,  immediately  reunite,  reproducing  the  original 
compound,  which  then  settles  or  condenses  on  the  cool  part  of  the 
tube,  forming  a  sublimate. 

[Generally,  however,  a  certain  small  amount  of  the  dissociated 
portions  of  the  compound  escapes  recombination  :  e.g.  heat  a  small 
quantity  of  ammonium  chloride  in  a  dry  test-tube  ;  notice  that 
white  fumes  are  produced,  which  sublime  up  the  tube.  Now  hold 
a  moistened  red  litmus  paper  in  the  mouth  of  the  tube,  and  it  will 
be  turned  blue,  showing  that  a  portion  of  the  ammonia  escapes 
from  the  tube  before  it  meets  the  hydrochloric  acid  from  which 
it  has  been  dissociated.  For  a  moment  discontinue  heating,  and 
presently  the  blued  paper  will  be  reddened,  for  the  molecules  of 
hydrochloric  acid  which  have  lost  their  partners  (the  escaped 
ammonia)  now  make  their  way  up  the  tube  and  act  on  the  litmus 
paper.] 


2O  Qualitative  Analysis. 

(b}  If  the  acid  is  non-  volatile,  or  volatile  only  at  a  high  tempera- 
ture, then  the  ammonium  salt  is  decomposed,  ammonia  being 
evolved,  while  the  acid  remains. 


.  Heat  a  little  ammonium  sulphate  or  phosphate  in  a  test- 
tube  ;  ammonia  is  rapidly  evolved,  and  may  be  detected  by  its 
characteristic  smell.] 

(c]  The  ammonium  salts  of  certain  oxyacids  which  readily  part 
with  oxygen  (such  as  ammonium  nitrate,  nitrite,  chromate)  are  also 
decomposed  by  heat,  the  ammonia  being  oxidised  to  nitrogen  or 
oxides  of  nitrogen. 

[E.g.  Heat  a  few  crystals  of  ammonium  nitrate  in  a  test-tube. 
Examine  the  gas  with  a  taper  and  a  glowing  splint  of  wood.] 

NH4NO3  =  2H2O  +  N2O 
NH4NO2  =  2H2O  +  N2 
(NH4)2Cr2O7  =  Cr2O3  +  4H2O  +  N2 

Ammonium  is  separated  from  the  other  members  of  the  group 
by  evaporating  the  solution  to  dryness,  and  strongly  heating  the 
residue  until  the  ammonia  is  completely  expelled,which  may  generally 
be  regarded  as  accomplished  when  fumes  are  no  longer  given  off. 

WET  REACTIONS.  —  Ammonium  salts  are  all  soluble  in  water, 
therefore  it  is  only  in  concentrated  solutions  that  any  precipitations 
with  reagents  can  be  formed.  Use  ammonium  chloride. 

Caustic  alkalies  (NaHO  or  KHO)  and  oxides  or  hydroxides 
of  metals  of  the  alkaline  earths  (e.g.  CaO,  Ba(HO)2),  when  heated 
with  an  ammonium  salt,  cause  the  evolution  of  ammonia  gas,  NH3. 

(NH4)2SO4  +  2NaHO  =  Na2SO4  +  2H2O  +  2NH3 
2NH4C1  +  CaO  =-  CaCl2  +  H2O  +  2NH3 

In  practice,  sodium  hydroxide  solution  is  added  either  to  the  solid 
salt  or  to  its  solution  in  water,  and  the  mixture  gently  warmed. 
The  evolved  ammonia  may  be  recognized  (i)  by  its  characteristic 
odour  if  present  in  sufficient  quantities  ;  (2)  by  its  power  of  restoring 
the  blue  colour  to  moist  reddened  litmus  paper,  or  of  turning  tur- 
meric paper  brown  ;  (3)  by  the  formation  of  white  fumes  when  a 
glass  rod  moistened  with  strong  hydrochloric  acid  is  held  in  the 
mouth  of  the  test-tube. 

[In  special  cases,  as  in  the  examination  of  natural  waters,  minute 
traces  of  ammonia  are  detected  by  the  use  of  Nessler's  solution  (a 
solution  of  potassium  mercuric  iodide  in  potash),  which  gives  either 
a  brown  precipitate  or  a  coloration,  according  to  the  amount  of 
ammonia  present,  2(HgI2,2KI)  +  3KHO  +  NH3  =  NHg2"I,H2O 
+  7KI  +  2H,O.] 


Group    V.  2 1 

Hydrogen  platinum  chloride*  (chloroplatinic acid]  H2PtCl6, 
precipitates  from  concentrated  solutions  a  yellow  crystalline  com- 
pound, ammonium  chloroplatinate  (or  ammonium platinic  chloride) 
(NH4)2PtCl6,  soluble  in  170  parts  of  water  at  10°  ;  insoluble  in 
alcohol  and  ether.  This  compound  is  distinguished  from  the 
similar  potassium  salt  in  that,  when  strongly  heated,  it  leaves  a 
residue  of  spongy  platinum  only. 

Tartaric  acid,  H2(C4H4O6),  or  hydrogen  sodium  tartrate, 
HNa(C4H4O6),  produces  in  concentrated  solutions  a  white  pre- 
cipitate of  hydrogen  ammonium  tartrate,  H(NH4)(C4H4O6).  Soluble 
in  water,  in  mineral  acids,  and  in  alkalies;  insoluble  in  alcohol. 
This  compound  is  distinguished  from  the  corresponding  potassium 
salt  by  the  fact  that  when  strongly  heated,  the  carbonaceous  residue 
is  without  any  alkaline  reaction. 

[When  tartaric  acid  is  used,  the  acid  previously  in  combination 
with  the  ammonia  is  liberated  by  the  double  decomposition,  thus — 

NH4C1  +  H2(C4H406)  =  H(NH4)(C4H406)  +  HC1 

And  as  the  precipitate  is  soluble  in  mineral  acids,  the  delicacy  of  the 
reaction  is  increased  by  employing  hydrogen  sodium  tartrate  as  the 
reagent,  in  which  case  no  free  acid  is  formed  in  the  reaction,  thus — 

NH4C1  +  HNa(C4H4O6)  =  H(NH4)(C4H4O6)  +  NaCl] 

Sodium,  Na. 

DRY  REACTION. — Sodium  compounds,  when  heated  upon  a 
platinum  wire  in  a  Bunsen  flame,  undergo  volatilisation,  and 
impart  to  the  flame  a  brilliant  golden  yellow  colour.  This  flame- 
reaction  is  the  most  characteristic  and  delicate  test  for  this  metal,  f 

WET  REACTIONS. — All  sodium  salts  are  soluble  ;  sodium  platino- 
chloride  is  soluble  in  water,  in  alcohol,  and  in  ether.  Hydrogen 
sodium  tartrate  also  is  freely  soluble  in  water.  Sodium  pyroanti- 
monate,t  however,  is  less  soluble  in  water  than  the  corresponding 
potassium  salt,  and  is  therefore  precipitated  by  the  addition  of  a 

*  By  long  habit  this  reagent  is  called  platinum  chloride:  it  may  be 
regarded  as  platinum  chloride  plus  two  molecules  of  HC1 ;  PtCl4,  2HC1.  In 

solution  it  yields  H  and  PtCl6  ions. 

f  When  the  light  emitted  by  heating  a  sodium  salt  in  the  Bunsen  flame  is 
examined  by  the  spectroscope  (see  p.  27),  it  is  found  to  be  monochromatic,  i.e. 
to  consist  of  one  colour  only,  namely,  pure  yellow  light.  Many  common 
coloured  materials,  such  as  indigo,  have  the  power  of  absorbing  yellow  rays, 
hence  if  the  sodium  flame  be  viewed  through^,  thin  stratum  of  such  a  coloured 
solution,  the  yellow  light  is  entirely  intercepted.  For  the  use  that  is  made  of 
this  property,  see  potassium  (p.  22). 

J  Formerly  misnamed  sodium  metantimonate. 


22  Qualitative  Analysis. 

strong  solution  of  potassium  pyroantimonate  to  a  strong  solution  of 
a  sodium  salt,  such  as  sodium  chloride,  thus  — 

H2K2Sb2O7  +  2NaCl  =  H2Na2Sb2O7  +  2KC1 
Potassium,  K. 

DRY  REACTION. — When  potassium  compounds  are  heated  upon 
a  platinum  wire  in  a  Bunsen  flame,  they  impart  to  the  flame  a  pale 
violet  or  lilac  colour.  This  delicate  tint,  however,  is  completely 
masked  by  the  intense  yellow  colour  which  the  presence  of  even 
minute  quantities  of  sodium  compounds  impart  to  the  flame. 

Introduce  a  fragment  of  potassium  nitrate  into  the  Bunsen 
flame  upon  a  loop  of  clean  platinum  wire  ;  *  notice  the  lilac  colour 
imparted  to  the  flame.  Now  look  at  the  flame  through  a  potassio- 
scope,\  and  observe  that  it  appears  a  brilliant  crimson-red  colour. 
Upon  another  wire  introduce  a  particle  of  sodium  chloride  into  the 
flame,  and  notice  that  when  this  is  examined  through  the  potassio- 
scope, the  intense  golden  yellow  light  is  absolutely  cut  off,  and  is 
invisible.  Now  touch  the  wire  containing  the  nitre  with  a  fragment 
of  sodium  chloride,  and  again  bring  it  into  the  flame.  The  yellow 
of  the  sodium  completely  overpowers  and  masks  the  violet  of  the 
potassium  when  viewed  direct,  but  if  looked  at  through  the  potassio- 
scope,  the  red  colour  due  to  the  potassium  shines  up  as  brilliantly  as 
before,  while  the  yellow  is  completely  intercepted  \  (see  also  p.  33). 

WET  REACTIONS. — Most  potassium  salts  are  soluble  in  water. 
Use  a  solution  of potassiiim  chloride. 

*  By  merely  touching  the  wire  with  the  fingers,  it  contracts  sufficient  sodium 
compounds  to  give  the  yellow  flame.     To  clean  it,  it  should  be  dipped  in  hydro- 
chloric acid  and  heated  until  it  ceases  to 
impart  any  colour  to  the  flame. 

f  The  potassioscope  consists  merely 
of  a  small,  flat,  glass  cell,  containing  a 
dilute  solution  of  one  of  the  aniline  blue 
dyes,  known  as  "soluble  blue  X.L." 
The  advantage  of  this  over  ordinary  blue 
glass  or  the  older  indigo  prism  lies  in  the 
fact  that  no  other  metal  but  potassium 
(except  the  extremely  rare  element  rubi- 
dium) gives  a  flame  which  appears  red 
when  viewed  through  the  potassioscope, 
whereas  lithium,  barium,  strontium,  and 
calcium  all  give  flames  which  appear  red 
through  indigo  or  blue  glass. 

\  When  studying  flame  reactions,  it  is 
often  of  the  greatest  convenience  to  use  a 
stand  on  which  to  support  the  platinum 
wire,  so  that  the  hands  may  be  free ;  a 
simple  stand  is  readily  constructed  as 
shown  in  Fig.  8.  A  piece  of  glass  tube 
or  glass  rod  is  inserted  in  a  large  cork 

pIG  8  (rubber,  being  heavier,  makes  a  steadier 

foot),  and  a  piece  of  galvanized  iron  wire 
is  twisted  two  or  three  times  round  the  rod  with  one  end  projecting  at  right 


Gro2tp   V.  23 


Hydrogen  platinum  chloride  (chloroplatinic  aci 
produces,  with  concentrated  solutions  of  potassium  salts,-  a  yellow 
crystalline  precipitate  of  potassium  chloro-platinate  (or  potassium 
platinic  chloride),  K2PtCl(;,  soluble  in  no  parts  of  water  at  10° 
(therefore  more  soluble  than  the  corresponding  ammonium  com- 
pound). Soluble  in  alkalies  (therefore  the  solutions  used  should 
be  acid)  ;  nearly  insoluble  in  alcohol  ;  quite  insoluble  in  a  mixture 
of  alcohol  and  ether  (therefore  the  precipitation  of  this  compound 
is  promoted  by  the  addition  of  alcohol). 

Hydrogen  sodium  tartrate,  HNa(C4H4O6),  gives,  with  solu- 
tions of  potassium  salts,  a  white  crystalline  precipitate  of  hydrogen 
potassium  tartrate,  HK(C4H4OG)  ;  soluble  in  much  water,  and  also 
in  acids  and  alkalies  (therefore  the  solution  should  be  both  concen- 
trated and  neutral).  The  precipitate  is  insoluble  in  alcohol. 

Hydrofluosilicic  acid  (or  silico-fluoric  acid),  H2SiFG, 
throws  down  a  white  precipitate  of  gelatinous  appearance,  consisting 
of  potassium  silicofluoride,  K2SiFfi,  sparingly  soluble  in  water. 

Magnesium,  Mg. 

In  the  "  natural  classification  "  of  the  elements,  magnesium  is 
associated  with  the  rnetals  of  the  alkaline  earths  (Be,  Ba,  Sr,  Ca) 
on  the  one  hand,  and  with  zinc  and  cadmium  on  the  other.  Its 
position  along  with  the  alkalies  in  Group  V.  of  the  analytical 
classification,  is  simply  because  it  differs  from  the  other  members  of 
its  own  natural  family  in  that  the  presence  of  ammonium  chloride 
prevents  the  precipitation  of  magnesium  hydroxide  by  ammonia 
in  Group  III.,  and  also  of  magnesium  carbonate  by  the  group 
reagent  of  Group  IV. 

DRY  REACTION.  —  When  magnesium  salts  are  strongly  heated 
in  the  outer  blowpipe,  a  white  infusible  residue  of  the  oxide  is  left. 
If,  after  cooling,  the  residue  be  moistened  with  a  drop  or  two  of 
cobalt  nitrate  solution  and  again  strongly  heated  in  the  outer  blow- 
pipe flame,  the  mass  acquires  a  pink  colour.  This  reaction  is 
reliable  only  in  the  absence  of  other  metallic  oxides. 

WET  REACTIONS.  —  Of  the  common  salts  of  magnesium,  the 
sulphate,  chromate,  nitrate,  and  halogen  salts  are  soluble  in  water. 
One  prominent  characteristic  of  magnesium  compounds  is  the 
readiness  with  which  they  form  "  double  "  salts,  many  of  which  are 
soluble  in  water.  Use  magnesium  sulphate. 

angles  to  the  upright.  The  little  glass  tube  into  which  the  platinum  wire  is 
fused,  is  then  slipped  over  the  projecting  iron  wire.  This  arrangement  admits 
of  the  wire  being  raised  or  lowered  as  desired,  while  at  the  same  time  it  readily 
remains  in  any  position. 


24  Qualitative  Analysis. 

Alkaline  hydroxides  (NH4HO,  KHO,  NaHO,  Ca(HO)2,  or 
Ba(HO)2)  precipitate  from  solutions  of  magnesium  sulphate  or 
chloride,  white  magnesium  hydroxide,  Mg(HO)2.  Almost  insoluble 
in  water  ;  soluble  in  acids,  soluble  in  ammonium  chloride — 

Mg(HO)2  +  4NH4C1  =  2NH4HO  +  MgCl2,2NH4Cl  (soluble 
double  salt) 

Owing  to  the  solubility  of  magnesium  hydroxide  in  ammonium 
chloride,  only  half  the  magnesium  is  precipitated  from  magnesium 
chloride  by  means  of  ammonia,  thus — 

2MgCl2  +  2NH4HO  =  Mg(HO)2  +  MgCl2,2NH4Cl 

If  ammonium  chloride  is  previously  present  in  sufficient  quantity 
the  alkaline  hydroxide  gives  no  precipitate. 

Stated  in  terms  of  the  ionic  theory,  the  solubility  of  magnesium 
hydroxide  in  ammonium  chloride  is  due  to  two  causes  ;  first  the 
slight  solubility  of  the  compound  in  water,  with  a  corresponding 

slight  dissociation  into  Mg  and  HO  ions  ;  and  second  the  com- 
paratively small  extent  to  which  ammonium  hydroxide  is  ionised. 
The  addition  of  ammonium  chloride  throws  into  the  solution  a  large 

excess  of  NH4  ions,  one  effect  of  which  is  to  reduce  the  ionisation 
of  ammonium  hydroxide  to  such  an  extent  that  there  are  no  HO 

ions  available  for  union  with  Mg  ions  ;  hence  no  precipitate  of 
Mg(HO)2  occurs.  On  the  other  hand,  if  ammonium  chloride  be 

added  after  the  precipitation  of  Mg(HO)2,  then  the  NH4  ions  thus 

introduced  unite  with  the  HO  ions  provided  by  the  slightly  ionised 
magnesium  hydroxide,  yielding  NH4HO  (practically  unionised). 

To  restore  the  equilibrium  disturbed  by  this  removal  of  HO  ions, 
more  of  the  Mg(HO)2  undergoes  ionisation,  and  this  process 
continues  until  the  whole  of  the  Mg(HO)2  has  passed  into  solution. 
The  above  equation  then  becomes — 

Mg,  2HO  +  4NH4,  4.C1  =  2NH4HO  +  [Mg,  Cl,  Cl,  2N+H4,  2C1] 

ionised  double  salt. 

Alkaline  carbonates  (K2CO3,  Na2CO3,  (NH4)2CO3)  pro- 
duce in  solutions  of  magnesium  salts,  in  the  absence  of  ammonium 
salts,  precipitates  of  basic  carbonates  of  magnesium,  the  composition 
of  which  varies  with  conditions  of  temperature  and  concentration. 
The  precipitate  with  (NH4)2CO3  only  separates  out  after  a  short 


Detection  of  the  Metals  of  Group   V.  25 

time.  In  the  presence  of  ammonium  chloride  these  reagents  give  no 
precipitate. 

Hydrogen  disodium  phosphate,  HNa2PO4,  precipitates 
hydrogen  magnesium  phosphate,  HMgPO4,  and  tri-magnesium 
phosphate,  Mg3(PO4)2.  In  the  presence  of  ammonium  chloride, 
however,  the  double  ammonium  magnesium  phosphate  is  thrown 
down  as  a  white  crystalline  precipitate,  NH4MgPO4.  It  is 
appreciably  soluble  in  water,  but  insoluble  in  ammonia;  hence 
ammonia  must  be  previously  added. 

In  very  dilute  solutions  the  precipitation  only  takes  place  on 
long  standing.  It  is  accelerated  by  stirring  with  a  glass  rod,  the 
deposition  first  appearing  where  the  rod  has  rubbed  the  glass 
vessel.  The  precipitate  is  soluble  in  acids,  even  acetic  acid,  but 
reprecipitated  by  ammonia. 

SYSTEMATIC  PLAN  OF  ANALYSIS  FOR  THE  METALS  OF 
GROUP  V. 

After  having  carefully  gone  through  the  various  reactions  for 
the  metals  of  Group  V.,  the  student  should  proceed  to  the  examina- 
tion of  a  few  solutions  containing  mixtures  of  two  or  more  salts  of 
these  metals. 

The  various  special  tests  by  which  the  individual  members  of 
this  group  are  recognised  are  not,  for  the  most  part,  interfered  with 
by  the  presence  of  the  rest  of  the  group.  Therefore  a  complete 
separation  of  all  the  metals  is  not  necessary.  Thus,  the  test  for 
ammonium  (evolution  of  ammonia  by  heating  with  sodium 
hydroxide)  can  be  made  in  the  presence  of  Mg,  K,  and  Na  com- 
pounds ;  and  obviously  must  be  made  in  a  separate  portion  of  the 
solution  under  examination  from  that  in  which  sodium  is  to  be 
tested  for,  as  it  involves  the  addition  of  a  sodium  compound. 

The  test  for  magnesium,  likewise  (precipitation  of  NH4MgPO4), 
may  be  made  in  the  presence  of  all  the  other  members  ;  and 
clearly  must  be  made  also  in  a  separate  portion  of  the  solution,  as 
it  involves  the  addition  of  both  ammonium  and  sodium  compounds. 
Similarly,  the  flame  tests  for  potassium  and  sodium  are  not  inter- 
fered with  by  the  presence  of  ammonium  or  magnesium.  The 
flame  reaction  for  potassium,  however,  unless  examined  by  the  aid 
of  the  potassioscope,  must  always  be  corroborated  by  the  forma- 
tion of  potassium  chloroplatinate.  But  before  this  test  can  be 
applied,  ammonium  salts  must  first  be  removed. 

Solutions  may  be  examined  for  the  metals  of  Group  V., 
NH4,  Na,  K,  Mg,  by  the  following  system  :— 


26  Qualitative  Analysis. 

DETECTION  OF  THE  METALS  OF  GROUP  V. 

Operation  i. — To  a  portion  of  the  solution  add  NaHO,  and  heat 
in  a  test-tube.  The  evolution  of  ammonia  (detected  by  its  odour, 
and  its  action  on  test-papers)  proves  the  presence  of  NH4. 

Operation  2. — To  a  second  portion  add  NH4C1,  NH4HO,  and 
HNa2PO4.  A  white  crystalline  precipitate  of  NH4MgPO4  proves 
the  presence  of  Mg. 

Operation  3. — Evaporate  another  (and  larger)  portion  to  dryness 
in  a  porcelain  dish.  If  ammonium  salts  are  present  *  (already 
ascertained  in  Operation  i),  they  must  be  removed.  For  this  pur- 
pose, scrape  the  residue  out  of  the  dish  and  strongly  heat  it  on  the 
lid  of  a  platinum  crucible  (or  a  piece  of  platinum  foil),  until,  on 
momentarily  withdrawing  it  from  the  flame,  fumes  are  no  longer 
visible. 

Dissolve  the  residue  in  a  small  quantity  of  water,  and  add  one 
drop  of  HC1.  Dip  a  clean  platinum  loop  into  the  solution,  and  heat 
it  in  a  Bunsen  flame.  An  intense  f  yellow  coloration  proves  the 
presence  of  Na.  A  lilac  colour  indicates  the  presence  of  K. 

In  either  case  examine  the  flame  through  the  potassioscope  ;  a 
crimson  flame  indicates  K. 

Add  to  the  solution  of  the  residue  a  few  drops  of  H2PtCl6,  and 
stir  with  a  glass  rod.  A  yellow  precipitate  of  K2PtCl6  confirms  K. 


APPENDIX    TO    CHAPTER    III 

LITHIUM,   RUBIDIUM,   AND   CAESIUM. 

These  three  elements  are  usually  placed  in  the  category  of  rare 
metals.  It  must  be  remembered,  however,  that  there  are  degrees  of 
rarity  ;  and  while  the  compounds  of  rubidium  and  cassium  are 
certainly  among  the  very  rare  substances  with  which  the  chemist 
comes  into  contact,  those  of  lithium,  on  the  other  hand,  are  very 
widely  distributed  and  are  much  more  frequently  met  with.J 

*  In  a  complete  analysis,  ammonium  salts  are  always  present  here,  as  they 
will  have  been  introduced  in  the  process  of  separating  the  other  groups. 
Under  these  circumstances,  therefore,  the  operation  of  removing  ammonium 
compound's  is  always  necessary.  The  substance  under  analysis  is  tested  for 
ammonium  before  the  ammoniacal  reagents  are  introduced. 

f  More  or  less  of  a  yellow  flame  is  usually  obtained,  owing  to  the  presence 
of  traces  of  sodium  compounds  as  impurities  in  the  reagents  previously  used  in 
separating  the  groups  in  the  course  of  a  complete  analysis. 

t  Perhaps  a  rough  idea  of  the  relative  rarity  of  the  compounds  of  these 
metals  might  be  gained  by  a  comparison  of  their  cost.  Lithium  salts  can  be 
obtained  for  about  i2s.  per  pound,  while  rubidium  and  caesium  salts  cost 
about  5-r.  per  drachm,  i.e.  at  the  rate  of  ^64  per  pound. 


The  Rare  Metals  of  Group   V  27 

Lithium,  Li. 

DRY  REACTION. — Lithium  salts  impart  to  the  flame  a  brilliant 
carmine-red  colour. 

WET  REACTIONS. — All  the  common  salts  are  readily  soluble  in 
water,  except  the  carbonate,  phosphate,  and  oxide,  which  are 
soluble  with  difficulty.  The  chloride  and  nitrate  are  soluble  in  a 
mixture  of  alcohol  and  ether  (distinction  from  Na  and  K,  the 
chlorides  and  nitrates  of  which  are  not  soluble). 

Na2CO3  and  K2C03  precipitate  Li2CO3  from  cold  moderately 
concentrated  solutions  (i  part  dissolves  in  100  parts  of  water). 

HNa2PO4  gives  a  white  precipitate,  on  boiling,  of  Li3PO4.  The 
precipitation  is  complete  in  the  presence  of  NaHO. 

PtCl4  gives  no  precipitate  (distinction  from  NH4,  K,  Rb,  Cs). 

HNa(C4H4O6),  hydrogen  sodium  tartrate,  gives  no  precipitate. 

Rubidium,  Kb,  and  Caesium,  Cs. 

The  compounds  of  these  metals  present  the  very  closest 
resemblance  to  those  of  potassium,  and  there  are  scarcely  any 
chemical  reactions  by  which  they  can  be  distinguished.  The 
separation  of  these  metals  is  based  on  the  different  degrees  of 
solubility  of  their  chloro-platinates.  In  this  respect,  as  well  as  in 
their  other  properties,*  rubidium  stands  intermediate  between 
potassium  and  caesium. 

When  heated  in  a  Bunsen  flame,  rubidium  and  caesium  salts 
impart  to  it  a  lilac  colour,  which  to  the  unaided  eye  is  absolutely 
indistinguishable  from  that  produced  by  potassium  compounds  ; 
and  when  compounds  of  these  metals,  as  well  as  those  of  lithium, 
are  mixed  with  comparatively  minute  quantities  of  sodium  salts,  the 
colours  they  give  to  the  flame  are  completely  overpowered  and 
masked  by  the  yellow  of  the  sodium.  By  means  of  the  spectroscope, 
however,  not  only  are  the  apparently  identical  colours  given  by 
potassium,  rubidium,  and  caesium  proved  to  consist  of  light  of 
different  quality  or  composition,  but  the  presence  of  any  or  all  of 
them  is  easily  and  certainly  detected  even  when  admixed  with 
sodium  salts. 

The  spectroscope  is  an  instrument  by  means  of  which  the  light 
emitted  by  strongly  heated  substances  can  be  examined  after  it  has 
been  made  to  pass  through  a  glass  prism.  Its  use  depends  upon 
the  fact  that  different  coloured  lights  possess  different  degrees  of 
refrangibility  ;  that  is  to  say,  different  coloured  rays  of  light  are 
bent  out  of  their  straight  course,  by  passage  through  a  prism,  at 
different  angles.  Ordinary  white  light  is  composed  of  rays  of  all 
degrees  of  refrangibility,  hence,  when  such  light  passes  through  a 

*  Atomic  weights,  melting-points,  etc.  Also  in  the  optical  properties  of 
their  crystallized  salts  (Tutton,  J.  C.  S.,  May,  1896). 


28 


Qualitative  Analysis, 


prism,  the  various  coloured  rays  are  separated,  and  spread  out  in  the 
order  of  their  refrangibility,  the  least  refrangible  red  at  one  extreme 
to  the  deep  violet  at  the  other.  This  familiar  " rainbow"  coloured 
band  of  light  is  called  the  continuous  spectrum. 

In  the  spectroscope  the  light  is  passed  through  a  narrow  slit  at 
one  end  of  a  small  telescope,  and  an  image  of  the  slit  is  received 
upon  a  glass  prism.  This  bends  the  light  out  of  its  straight  course, 
and  spreads  it  out  into  the  various  colours  of  which  it  is  composed' 
If  white  light  be  admitted,  then  the  continuous  spectrum  is  seen, 


FIG.  9. 

which  is  an  infinite  number  of  images  of  the  slit  arranged  side  by 
side  ;  if  such  a  w0#0chromatic  light  as  that  given  by  heating  sodium 
salts  in  the  flame  be  used,  then  one  image  of  the  slit  is  seen  in  that 
part  of  the  spectrum  which  corresponds  to  the  particular  degree  of 
refrangibility  of  the  light.  In  this  case  we  say  that  the  spectrum 
of  sodium  consists  of  one  line  (that  is,  one  image  of  the  slit)  in  the 
yellow,  or  one  yellow  line*  The  light  which  passes  out  of  the  prism 
is  usually  examined  by  a  second  telescope,  which  can  be  revolved 
round  the  prism  so  that  it  can  sweep  the  whole  length  of  the 
spectrum. 

When    the   red   colour   imparted    to   the    Bunsen   flame   by  a 

*  When  this  sodium  line  is  examined  by  a  higher  dispersive  power,  it  is 
found  to  consist  of  a  group  of  lines. 


The  Rare  Metals  of  Group   V.  29 

lithium  salt  is  examined  by  the  prism,  it  is  seen  to  consist  of  light 
of  two  degrees  of  refrangibility,  therefore  two  images  of  the  slit  are 
seen,  one  in  the  bright  orange  (a  little  on  the  red  side  of  the  sodium 
line),  and  the  other  an  extremely  brilliant  image  in  the  red.  The 
spectrum  of  lithium,  therefore,  is  one  bright  orange  line  and  one 
brilliant  red  line  (see  Fig.  9). 

Similarly,  when  the  lilac  flame  of  potassium  is  examined  in  this 
way,  the  light  is  found  to  be  composed  of  three  colours  ;  therefore 
three  lines  are  seen — one  extremely  bright  line  far  down  in  the  red  ; 
a  second,  less  brilliant,  about  halfway  between  the  first  and  the 
yellow  sodium  (or  D)  line  ;  while  the  third  is  far  away  in  the  deep 
violet,  and  not  easy  to  see.* 

If  a  mixture  of  salts  of  sodium,  potassium,  and  lithium  be  heated 
in  a  flame,  and  the  light  examined  by  the  spectroscope,!  then  all 
six  lines  are  seen  ;  sodium,  one  yellow ;  potassium,  two  red  and  one 
violet;  lithium,  one  red  and  one  orange. 

In  such  a  simple  mixture  as  this  there  is  no  difficulty  in 
identifying  the  various  lines  of  the  different  elements,  but  in  more 
complex  spectra  it  is  only  possible  to  do  so  by  recording  the  exact 
position  they  each  occupy  on  some  fixed  scale. 

When  the  spectra  of  rubidium  and  caesium  are  in  this  way  com- 
pared with  that  of  potassium,  a  striking  difference  is  at  once 
observed.  In  the  first  place,  they  are  obviously  much  more  complex. 
Rubidium  shows  two  strongly  marked  red  lines  (hence  the  name 
of  the  element)  very  close  together,  and  near  to  the  red  line  of 
potassium.  Then  a  number  of  lines  in  the  orange-red  and  in  the 
green,  and  lastly  two  characteristic  and  brilliant  violet  lines. 

Caesium  gives  red  lines  (not  so  far  down  as  those  of  rubidium) 
orange  and  green  lines,  and  two  most  characteristic  brilliant 
lines  in  the  bright  blue  part  of  the  spectrum  (ccesium  signifies 
sky -blue}. 

In  the  chart  (Fig.  9),  the  most  prominent  lines  f  in  the  spectra 
of  the  five  metals  of  the  alkalies  are  compared  together.  The 
letters  at  the  top  refer  to  lines  in  the  solar  spectrum  mapped  by 
Fraunhofer,  which  correspond  to  those  given  by  known  metals. 
Thus,  the  D  line  in  the  solar  spectrum  corresponds  to  the  yellow 
line  of  sodium. 

*  The  sodium  D  line  will  be  also  visible,  not  only  on  account  of  the  presence 
of  traces  of  sodium  compounds  in  the  other  salts,  but  owing  also  to  the 
presence  of  particles  of  salt  which  are  always  floating  about  in  the  air,  and 
which  tinge  the  Bunsen  flame. 

t  Where  the  equipment  of  the  laboratory  will  allow,  every  student  should 
have  the  opportunity  of  examining  the  spectra  of  the  metals  of  the  alkalies  and 
alkaline  earths. 

£  Speaking  generally,  it  may  be  said  that  at  higher  temperatures  than  can  be 
obtained  in  the  Bunsen  flame,  lines  will  appear  in  the  spectrum  of  a  particular 
metal  which  are  not  seen  at  the  lower  temperature. 


CHAPTER   IV. 
REACTIONS  OF  THE  METALS  OF  GROUP  IV. 

THIS  group  consists  of  the  three  metals,  barium,  strontium, 
and  calcium,  known  as  the  metals  of  the  alkaline  earths. 

The  group-reagent,  ammonium  carbonate,  precipitates  the 
carbonates  of  the  metals  in  an  ammoniacal  solution,  even  in  the 
presence  of  ammonium  chloride.  Hence  the  group-reagent  separates 
these  three  metals  from  magnesium. 

Barium,  Ba. 

DRY  REACTION. — Barium  compounds,  heated  on  platinum  wire 
in  the  Bunsen  flame,  impart  a  pale  apple-green  colour  to  the  flame, 
which  becomes  more  distinct  if  the  substance  on  the  wire  is 
moistened  with  strong  hydrochloric  acid.  The  test  is  not  very 
reliable. 

Barium  sulphate,  BaS04  (also  SrSO4  and  CaSO4),  when  heated 
on  charcoal  or  with  carbon,  is  reduced  to  the  sulphide. 

WET  REACTIONS.— Of  the  common  salts  of  barium,  the  chloride, 
bromide,  iodide,  nitrate,  chlorate,  acetate,  and  sulphide  are  soluble 
in  water. 

Ammonium  carbonate,  (NH4)2CO3,  group-reagent  (also 
Na2CO3  and  K2CO3)  precipitates  barium  carbonate,  BaCO3,  as  a 
white  amorphous  powder.  Insoluble  in  water  ;  readily  dissolved, 
with  evolution  of  carbon  dioxide,  by  dilute  acids ;  slightly  soluble  in 
NH4C1. 

It  is  also  dissolved  by  a  solution  of  carbon  dioxide  in  water, 
forming  hydrogen  barium  carbonate,  H2Ba(CO3)2.* 

*  On  this  account  complete  precipitation  of  barium  carbonate  cannot  be 
accomplished  by  hydrogen  ammonium  carbonate,  HNH4CO3,  for  carbon 
dioxide  is  evolved,  which  dissolves  a  portion  of  the  precipitate — 

2HNH4C03  +  BaCl2  =  BaCO3  +  2NH4C1  +  HaO  f  CO2 

And  since  the  solution  of  ammonium  carbonate  used  as  the  group-reagent 
consists  chiefly  of  HNH4CO3  (because  the  normal  salt  undergoes  decomposition 
into  the  hydrogen  salt  and  free  ammonia  when  in  aqueous  solution),  it  is  always 
necessary  to  first  add  NH4HO  before  applying  the  group-reagent. 


Group  IV.  31 

H  SO,,  or  any  soluble  sulphate,  produces  a  white  granular 
precipitate  of  BaSO4,  practically  insoluble  in  water,  insoluble  also 
in  acids  and  alkalies.  (Boiling  concentrated  H2SO4  slowly  dis- 
solves it,  forming  hydrogen  barium  sulphate,  H2Ba(SO4)2.)  In- 
soluble in  solutions  of  (NH4)2SO4.  BaSO4,  being  practically  insoluble 
in  water,  is  precipitated  by  a  saturated  solution  of  SrSO4,  although 
such  a  solution  contains  only  I  part  of  salt  in  7000  parts  of  water. 

Potassium  chromate,  K2CrO4,  produces  a  primrose-yellow 
precipitate  of  barium  chromate,  BaCrO4,  practically  insoluble  in 
water  (distinction  from  SrCrO4,  which  is  MODERATELY  soluble,  ana 
CaCrO4,  which  is  VERY  soluble  in  water].  It  is  insoluble  in  acetic 
acid  (distinction  from  SrCrO4) ;  soluble  in  HNO3  and  in  HC1. 
By  boiling  with  K2CO3,  barium  chromate  is  converted  into  the 
carbonate— 

BaCr04  +  K2CO3  =  BaCO3  +  K2CrO4 

Hydrofluo  silicic  acid,  H2SiF6,  gives  a  white  crystalline  pre- 
cipitate of  barium  siiicofluoride,  BaSiF6,  slightly  soluble  in  water, 
but  insoluble  on  the  addition  of  alcohol  (distinction  from  SrSiF6 
and  CaSiF6,  which  are  readily  soluble  in  water). 

Strontium,  Sr. 

DRY  REACTION. — When  heated  in  the  Bunsen  flame,  volatile 
strontium  salts,  such  as  SrCl2,  Sr(NO3)2,  impart  a  rich  crimson 
colour  to  the  flame  ;  other  salts  require  to  be  moistened  upon  the 
wire  with  strong  HC1.  Strontium  sulphate  must  be  converted  into 
the  sulphide,  which  is  then  dissolved  in  HC1  and  the  chloride 
tested  in  the  flame.* 

WET  REACTIONS.— The  same  common  salts  of  strontium  as  ot 
barium  are  soluble  in  water.  The  chromate  and  sulphate  are 
somewhat  soluble. 

(NH4)2CO3  (Na2CO3  and  K2CO3)  precipitates  SrCO3,  exactly 
similar  to  the  barium  compound  in  its  reactions. 

H2SO4  or  soluble  sulphates  precipitate  SrSO4.  The  precipitate 
is  slightly  soluble  in  water  (i  :  7000),  considerably  soluble  in  HC1 
or  HNO3,  but  almost  insoluble  in  a  solution  of  (NH4)2SO4  (dis- 
tinction from  CaSO4).  SrSO4  is  precipitated  by  a  solution  of 

*  Small  quantities  of  SrSO4  upon  a  filter  may  be  converted  into  SrS  by  first 
drying  the  filter,  then  folding  it  into  a  small  roll,  and  twisting  a  platinum  wire 
round  it  so  as  to  hold  it  in  a  flame.  The  paper  is  burnt  to  an  ash,  and  the 
latter  is  then  heated  upon  the  wire  in  an  ordinary  smoky  gas-flame,  when  the 
reduction  of  the  sulphate  takes  place.  The  ash  is  then  touched  with  a  glass 
rod  dipped  in  strong  HC1,  and  on  bringing  it  into  a  Bunsen  flame  the  crimson 
flame  is  seen. 


32  Qualitative  Analysis. 

CaSO4  *  (a  saturated  solution  of  which  only  contains  i  part  in 
430  parts  of  water)  ;  the  precipitation  does  not  take  place  at  once 
in  cold  solutions,  but  appears  quickly  on  heating. 

K,CrO4  and  H2SiF6  (see  barium  reactions  with  these  reagents) 

Calcium,  Ca. 

D  RY  REACTIONS, — Calcium  compounds,  when  heated  in  a  Bunsen 
flame,  impart  to  it  a  reddish  colour,  especially  if  previously  moistened 
with  hydrochloric  acid.  The  presence  of  strontium  masks  the  red 
colour  given  by  calcium  compounds.  When  calcium  carbonate  is 
strongly  heated  it  loses  carbon  dioxide,  and  is  converted  into 
calcium  oxide — 

CaCO3  =  CO2  +  CaO 

This  process  is  carried  on  in  the  lime-kilns  during  the  operation  of 
"  burning  lime."  Limestone  (calcium  carbonate)  is  thus  converted 
into  lime  (calcium  oxide],  SrCO3  and  BaCO3  undergo  similar 
decomposition ;  SrCO3  less  readily  than  CaCO3,  while  BaCO3 
requires  prolonged  heating  to  a  white  heat  to  effect  the  change. 

WET  REACTIONS. — The  same  common  salts  of  calcium  are 
soluble  in  water  as  of  strontium  and  barium  ;  and,  generally 
speaking,  calcium  salts  are  more  readily  soluble.  Thus,  the  chloride 
and  nitrate  are  extremely  deliquescent ;  the  sulphate  dissolves  to 
the  extent  of  i  :  430  (compare  Sr  and  Ba).  The  solubility  of  the 
oxalates,  however,  is  in  the  reverse  order,  calcium  oxalate  being 
the  most  insoluble,  and  barium  oxalate  the  most  soluble.  A 
peculiarity  of  calcium  salts  is  that  many  of  them  appear  to  be  less 
soluble  in  hot  than  in  cold  water. 

(NH4)2CO3  (also  Na2CO3  and  K2CO3)  precipitates  CaCO3, 
similar  to  the  barium  and  strontium  compounds  in  its  reactions. 

H  SO,  or  soluble  sulphates,f  added  to  a  strong  solution  of  a 
calcium  salt,t  give  an  immediate  precipitate  of  calcium  sulphate. 

*  SrSO4  is  itself  soluble  in  water  to  the  extent  of  i :  7000,  but  obviously  the 
addition  of  a  solution  of  CaSO4  or  HijSO4  to  this  solution  does  not  precipitate 
the  SrSO4.  Neither  is  the  strontium  precipitated  from  this  solution  by  any 
reagent  which  forms  with  it  a  compound  which  is  more  soluble  in  water  than 
is  the  sulphate,  e.g.  oxalic  acid  or  oxalates. 

f  The  soluble  sulphate  must  be  more  soluble  than  calcium  sulphate  ;  SrSO4 
therefore,  being  less  soluble,  would  not  give  a  precipitate  with  a  solution  of  a 
calcium  salt,  although,  being  more  soluble  than  BaSO4,  it  throws  down  this 
compound  from  solutions  of  barium  salts.  For  the  same  reason,  a  solution  of 
CaSO4  will  throw  down  SrSO4  and  BaSO4  from  solutions  of  strontium  and 
barium  salts. 

%  Only  in  the  case  of  calcium  salts  which  arc  more  soluble  than  calcium 
sulphate. 


Group  IV.  33 

From  more  dilute  solutions  the  precipitate  only  separates  after 
some  time,  or,  if  still  more  dilute,  not  at  all.  The  precipitate  is 
insoluble  in  alcohol,  therefore  the  addition  of  this  liquid  in 
considerable  bulk  favours  the  precipitation. 

Calcium  sulphate  is  readily  soluble  in  a  concentrated  solution 
of  ammonium  sulphate,  especially  when  hot  (distinction  from 
SrSO4  and  BaSO4).  Boiling  with  potassium  carbonate  easily 
converts  it  into  calcium  carbonate. 

Oxalic  acid,  H2C2O4,  or  soluble  oxalates,  gives  a  white 
crystalline  precipitate  of  calcium  oxalate.  The  precipitate  is 
soluble  in  mineral  acids  ;  hence,  when  oxalic  acid  is  used  as  the 
precipitant,  precipitation  is  not  complete,  owing  to  the  liberation 
of  the  mineral  acid  of  the  calcium  salts,  thus— 

CaCl2  +  H2C204  =  CaC204  +  2HC1 

Calcium  oxalate  is  insoluble  in  NH4HO,  therefore,  if  the  solution 
be  first  rendered  alkaline,  the  whole  of  the  oxalate  is  thrown  down. 
In  practice  it  is  usual  to  employ  ammonium  oxalate,  (NH4)2C2O4. 
In  cold  dilute  solutions  precipitation  is  not  immediate.  The  pre- 
cipitate is  insoluble  in  water  and  in  acetic  acid.  Strontium  and 
barium  oxalates  are  soluble,  to  a  small  extent,  both  in  water  and 
in  acetic  acid. 

SPECTRA  OF  BARIUM,  STRONTIUM,  AND  CALCIUM. 

The  spectra  given  by  these  metals,  when  their  salts  are  heated 
in  the  Bunsen  flame,  are  extremely  characteristic.  For  calcium, 
the  chloride  may  be  used,  with  occasional  moistening  with  hydro- 
chloric acid.  For  barium  and  strontium,  the  nitrates,  or,  better 
still,  the  chlorates  give  the  best  result,  although  in  the  absence  of 
these  salts  the  chlorides  can  be  employed. 

From  the  charts  of  their  spectra  given  in  Fig.  10,  it  will  be  seen 
that  they  each  give  lines  in  the  red  portion  of  the  spectrum.  Owing 
to  this  fact,  if  salts  of  these  metals  are  heated  in  a  Bunsen  flame, 
and  the  flame  examined  through  an  indigo  prism,  the  flame  will 
appear  ra/in  each  case,  because  indigo  does  not  intercept  or  absorb 
the  red  portion  of  the  spectrum.  The  potassioscope,  on  the  other 
hand,  has  the  property  of  cutting  off  the  red  rays  emitted  by  these 
metals.  It  absorbs  all  the  red  part  of  the  spectrum  down  very 
nearly  to  the  prominent  red  line  (K.o)  of  potassium  (compare  the 
chart  on  p.  28),  beyond  the  lithium  and  caesium  red  lines.  Hence, 
if  flames  which  are  coloured  by  the  alkaline  earths  are  examined 
through  this  instrument,  no  red  will  be  visible. 

D 


34  Qualitative  Analysis. 

The  importance  of  this  lies  in  the  following  fact :  When  Groups 
IV.  and  V.  are  separated  by  precipitating  the  carbonates  of  barium, 
strontium,  and  calcium,  minute  traces  of  these  metals  pass  through 
into  Group  V.,  as  their  carbonates  are  not  absolutely  insoluble 
Hence,  when  the  final  residue  is  obtained  (in  which  the  tests  for 


FIG.  10. 

potassium  are  made),  it  will  contain  these  traces  of  these  alkaline 
earths.  If  this  residue  be  tested  in  the  flame,  and  the  flame 
examined  by  the  potassioscope,  the  presence  of  these  members  of 
Group  IV.  will  not  vitiate  the  reaction  ;  but  with  the  indigo  prism 
the  red  light  they  emit  would  be  mistaken  for,  and  returned  as  that 
given  by  potassium.* 

SEPARATION  OF  GROUPS  IV.  AND  V. 

The  solution  is  first  rendered  alkaline  by  the  addition  of  NH4HO. 
NH4C1  is  then  added  (to  prevent  the  precipitation  of  magnesium 
carbonate),  after  which  ammonium  carbonate  is  added  until  the 
carbonates  of  the  metals  of  Group  IV.  are  completely  thrown  down. 
The  mixture  may  be  gently  warmed.  [It  must  not  be  boiled,  or 
the  precipitated  carbonates  will  react  with  the  NH4C1,  forming 
soluble  chlorides,  while  NH3  and  G02  will  escape  with  the  steam  ; 
thus,  BaC03  +  2NH4C1  =  BaCl2  +  CO2  +  2NH3  +  H2O.]  The 
mixture  is  filtered.  The  filtrate  is  examined  for  the  metals  of 
Group  V.  by  the  method  given  on  p.  25,  while  the  precipitate  is 
treated  in  the  following  way  : — 

*  The  only  element  which  could  be  mistaken  for  potassium,  when  using  the 
potassioscope,  is  the  exceedingly  rare  metal  rubidium,  whose  spectrum  contains 
red  lines  still  lower  down  than  those  of  potassium. 


Group  IV. 
SEPARATION  OF  THE  METALS  OF  GROUP  IV. 


35 


The  precipitate,  consisting  of  PaCO3,  SrCO3,  ai.d  CaCO3,  is  washed, 
and  then  dissolved  in  a  small  quantity  of  warm  dilute  acetic  acid, 
H(C  H3O2).  To  the  solution  thus  obtained  potassium  chromate, 
K2CrO4,  is  added  ;  the  mixture  gently  warmed  and  filtered 


I 

The  precipitate  con- 
sistsofyellow  barium 
chromate,  BaCrO4.  | 


The  filtratef  contains  strontium  and  calcium 
acetates.  Add  a  strong  solution  of 
(NH4)2SO4,  and  boil  for  a  short  time  to 
ensure" the  solution  of  CaSO4.  Filter. 


Boil  the  precipitate  with 

K2CO3,  which  con  verts 

BaCrO,  into  BaCO3. 
Filter,  and  wash    the 
precipitate   free  from 

The  precipitate  con- 
sists of  SrSO4. 

The  solution  contains 
calcium  sulphate. 
Add  ammonium  oxa- 

the  K2CrO4  .  Dissolve 
the  precipitate  in  HO, 
and  confirm  by  SrSCu 
or  H2SiF6. 

Wash  the  precipitate 
thoroughly,  and  con- 
firm by  the  flame 

late,  (NH4)2CoO4  (or 
(NH4)HO  and  oxa- 
lic acid,  H2C2O4). 

reaction. 

A  white  precipitate 

of  CaC2O4  confirms 

calcium. 

Other  methods  of  separating  the  metals  of  Group  IV.,  based  on 
the  insolubility  of  certain  of  their  common  salts  in  alcohol,  may 
be  employed.  Thus,  the  nitrates  of  barium  and  of  strontium  are 
insoluble  in  alcohol,  while  that  of  calcium  is  soluble  in  this  liquid 
(or  mixtures  of  alcohol  and  ether). 

The  precipitated  carbonates  are  dissolved  in  dilute  HNO3,  and 
the  solution  evaporated  to  dryness.  The  residue,  consisting  of  the 
mixed  nitrates,  is  then  treated  with  a  mixture  of  alcohol  and  ether, 
in  about  equal  volumes,  which  dissolves  the  Ca(NO3)2,  leaving 
Ba(N03)2  and  Sr(NO3)2. 

Again,  barium  chloride  is  insoluble  in  the  same  mixture,  whereas 
the  chlorides  of  strontium  and  calcium  are  both  soluble.  Therefore, 
if  the  mixed  carbonates  are  dissolved  in  dilute  HC1,  and  the  solu- 
tion evaporated  to  dryness,  the  SrCl2  and  CaCl2  may  be  dissolved 
out,  leaving  a  residue  of  BaCl2. 

*  Instead  of  adding  the  chromate  at  once  to  the  whole  of  the  solution,  a 
small  separate  portion  may  be  used,  and  the  presence  or  absence  of  the  Ba 
thus  ascertained.  If  present,  then  the  whole  of  the  solution  is  treated  as  above, 
and  the  precipitation  of  barium  chromate  may  then  be  taken  as  confirmatory. 
If  absent,  then  a  second  small  portion  of  the  solution  is  taken,  and  tested  for  Sr 
by  adding  CaSO4.  If  Sr  is  present,  then  the  separation  of  Sr  and  Ca  is  carried 
out  as  shown  above ;  but  if  CaSO4  gives  no  precipitate  even  on  warming, 
thus  proving  the  absence  of  Sr,  the  remainder  of  the  solution  is  at  once  tested 
for  Ca  by  the  oxalate  reaction. 

f  If  barium  has  been  separated  as  chromate,  the  filtrate  will  be  coloured 
yellow  by  the  excess  of  K2CrO4  used,  and  a  white  precipitate  may  appear 
yellow.  The  process  is  more  thorough  if  the  carbonates  of  Sr  and  Ca  are 
again  thrown  down  by  adding  ammonium  carbonate.  The  precipitate  is  then 
washed  free  from  chromate,  dissolved  in  a  few  .drops  ,of  HC1,  and  the  sulphates 
separated  as  directed  above. 


CHAPTER  V. 
THE  METALS  OF  GROUP  III. 

THIS  group  comprises  the  metals  aluminium,  chromium,  iron, 
manganese,  zinc,  nickel,  and  cobalt,  besides  a  number  of 
the  rare  metals  (see  Classification,  p.  16). 

Although  all  these  metals  are  precipitated  by  the  group-reagent, 
namely,  ammonium  sulphide,  in  the  presence  of  ammonia,  the 
compounds  that  are  thrown  down  are  not  all  of  similar  composition, 
as  was  the  case  in  Group  IV. 

The  addition  of  ammonia  to  a  solution  containing  the  metals 
of  this  group,  results  in  the  precipitation  of  the  hydroxides  (or 
hydrated  oxides)  of  Al,  Cr,  and  Fe,  corresponding  to  the  oxides 
R2O3.*  Thus,  taking  aluminium  in  potash  alum  as  an  example — 

K2S04,A12(S04)3  +  6NH4HO  -  K2SO4  +  3(NH4)2SO4  +  A12(HO)6 

In  the  presence  of  ammonium  chloride,  the  hydroxides  of  the 
metals  Mn,t  Zn,  Ni,  and  Co  (which  have  the  general  formula 
R(HO)2)  are  not  precipitated  by  ammonia,  for  the  same  reasons  as 
apply  in  the  case  of  magnesium.  (See  p.  24.) 

If  ammonium  sulphide  be  added  to  the  mixture  after  the 
ammonia,  the  already  precipitated  hydroxides  of  aluminium  and 
chromium  are  unchanged,  but  the  ferric  hydroxide  is  converted 
into  ferrous  sulphide,  with  the  elimination  of  sulphur  ;  thus — 

Fe2(HO)6  +  3H2S  =  2FeS  +  6H2O  +  S 
and  the  remaining   metals   are   also  thrown  down  as  sulphides. 

*  The  separation  of  the  metals  of  Group  III.  is  complicated  by  the  fact  that 
if  Mn,  Zn,  Ni,  and  Co  are  present  in  the  form  of  phosphates,  the  addition  of 
ammonia  causes  the  partial  precipitation  of  these  phosphates.  And,  further, 
if  the  members  of  Group  IV.  and  Mg  are  present  as  phosphates,  they  also  are 
thrown  down  by  ammonia,  and  therefore  appear  in  Group  III.  The  special 
treatment  of  phosphates  will  be  discussed  later. 

f  In  the  case  of  manganese,  the  separation  is  not  complete  under  all 
conditions  (see  reactions  of  manganese). 


Group  III. — Division  A.  37 

The  result,  therefore,  of  adding  NH4C1,  NH4HO,  and  (NH4)2S   is 
as  follows  : — 

Precipitated  by 
NH4HOinthe 


of|Fe2(HOy6  converted  into  FeS    * 


MnS 
ZnS 
NiS 
CoS 


Precipitated  on  the 
addition  of  (NH4)2S 


The  metals  of  Group  III.V  therefore,  may  be  subdivided  into 
two  families,  based  upon  their  behaviour  towards  ammonia  :  Divi- 
sion A,  consisting  of  the  three  metals,  aluminium,  chromium,  and 
iron ;  and  Division  B,  of  manganese,  zinc,  nickel,  and  cobalt. 

REACTIONS  o£  THE  METALS  OF  GROUP  III. — DIVISION  A. 
Aluminium,  Al. 

DRY  REACTION. — When  aluminium  compounds  are  strongly 
heated  on  charcoal  in  the  outer  flame,  aluminium  oxide  is  formed, 
and  if  this  be  moistened  with  a  solution  of  cobalt  nitrate,  and  again 
strongly  heated,  either  upon  the  charcoal  or  upon  a  loop  of  plati- 
num wire,  the  mass  assumes  a  rich  blue  colour,  due  to  the  formation 
of  cobalt  aluminate. 

This  test  is,  however,  greatly  masked  if  other  metallic  oxides 
which  are  coloured  are  present  at  the  same  time.  It  may  be 
employed  as  a  confirmatory  test  when  aluminium  is  separated  from 
iron  and  chromium  in  the  course  of  analysis. 

WET  REACTIONS. — Of  the  common  salts  of  aluminium,  the 
chloride,  A12C16,  and  sulphate,  A12(SO4)3,  are  soluble  in  water. 
The  important  salts,  however,  are  the  double  sulphates  of 
aluminium  with  ammonium  or  potassium,  known  as  ammo- 
nium alum,  (NH4)2SO4,A12(SO4)3,24H2O,  and  potassium  alum, 
K2SO4,A12(SO4)3,24H2O,  respectively.!  A  solution  of  either  of 
these  alums  may  be  used  for  the  following  reactions. 

*  In  reality  the  compounds  precipitated  are  the  hydrated  sulphides,  ex- 
pressed by  the  general  formula  R(HS)(HO),  or  RS,H2O.  To  avoid  unnecessary 
complication  in  reactions,  the  molecule  of  H2O  may  be  left  out  of  consideration. 

f  The  alums  constitute  a  large  class  of  double  sulphates,  having  the  general 
formula  M2SO4,R2(SO4)3)24H2O(  where  M  is  a  monovalent  element  or  group, 
such  as  potassium  or  ammonium,  and  R  is  either  aluminium,  iron,  chromium, 
or  manganese.  The  commonest  of  all  the  salts  is  potassium  aluminium  alum, 
K2Sp4,Al2(SO4)3,24H2O ;  this  is  the  salt,  therefore,  that  is  distinguished  by 
the  single  word  ALUM.  From  their  formulae  it  may  be  seen  at  a  glance  that 


3  8  Qualitative  Analysis. 

NH  HO  throws  down  a  white  translucent  precipitate  of  the 
hydrated  oxide,  or  hydroxide,*  ALO3,  3H2O,  or  A12(HO)6.  Soluble 
in  a  large  excess  of  the  reagent,  but  on  gently  boiling,  the  hydroxide 
in  entirely  precipitated.  [Prolonged  boiling,  however,  causes  partial 
dissociation  of  the  ammonium  salt  in  solution  into  ammonia,  which 
escapes,  and  free  acid,  which  then  begins  to  dissolve  the  preci- 
pitate.] In  the  presence  of  ammonium  chloride,  the  precipitation 
of  A12(HO)6  by  ammonia  is  complete.  The  precipitate  is  readily 
soluble  in  mineral  acids,  and  in  acetic  acid. 

KHO  or  NaHO  produces  the  same  precipitate,  readily  soluble 
in  an  excess  of  the  reagent,  forming  potassium  or  sodium  aluminate 
(Al2O3,3Na2O  or  Na6Al2O,;).  The  ready  formation  of  these 
aluminates  is  due  to  the  fact  that  aluminium  hydroxide  can  exhibit 
feeble  acidic  properties  ;  that  is  to  say,  it  is  ionised  to  some  extent 

into  H  and  A1O3  ions.     With  strong  bases,  therefore,  we  have  the 
reaction 

H6A12O6  +  6NaHO  =  6H2O  +  Na6Al2O6  or 


In  the  case  of  the  slightly  ionised  base  ammonium  hydroxide  it  is 
probable  that  ammonium  aluminate  is  first  formed,  and  then 
undergoes  hydrolysis,  this  latter  change  being  aided  by  the  gentle 
warming  which  determines  the  complete  precipitation  of  the 
hydroxide.  These  aluminates  are  decomposed  by  acids,  even  by 
such  feeble  acids  as  carbonic  acid  or  hydrosulphuric  acid 
(sulphuretted  hydrogen),  with  re-precipitation  of  the  aluminium 
hydroxide.  Thus,  with  carbonic  acid  — 


In  the  case  of  stronger  acids,  such  as  HC1,  the  same  action 
takes  place,  but  any  excess  of  the  acid  beyond  that  required  to 

these  salts  are  composed  of  a  molecule  of  each  of  the  two  sulphates,  together 
with  twenty-four  molecules  of  water.  With  compounds  of  this  description 
there,  is  unfortunately,  a  tendency  in  certain  quarters  to  add  the  formulas  of 
the  two  salts  together,  and  then  to  divide  the  numerals  by  their  greatest  com- 
mon measure  ;  thus,  KjSO4lAl,(SO4),l  24^0  =  K2A12(SO4)4,24H2O,  which, 
divided  by  2=  KAl)SO4)2li2H2O.  Presumably  this  plan  is  adopted  with 
a  view  to  simplification,  but  as  it  obscures  the  origin  and  the  nature  of  the 
compounds,  and  as  there  is  not  the  smallest  evidence  that  such  formulae  are 
more  exact  representations  of  the  molecular  constitution  of  the  compounds, 
their  use  is  greatly  to  be  deprecated. 

*  Three  hydrated  oxides  of  aluminium  are  known,  obtainable  by  precipi- 
tation under  different  circumstances,  namely,  AloO3,3H2O.  A12O3,2H2O, 
and  A12O3,H2O;  these  may  also  be  formulated  A12(HO)6,  A12O(HO)4,  and 
A12O^(HO)2  respectively.  The  first  of  these,  A12(HO)6  is  sometimes  written 
A1(HO)3.  (See  previous  note.) 


Group  III  —Division  A.  39 

combine  with  the  sodium  of  the  aluminate  at  once  re-dissolves  the 
A12(HO)6.  When,  therefore,  this  acid  is  used,  a  slight  excess  is 
added,  and  the  aluminium  hydroxide  is  re-precipitated  by  means 
of  ammonia. 

Sodium  and  potassium  aluminates  are  also  decomposed  by  am- 
monium chloride,  with  the  precipitation  of  aluminium  hydroxide  ;  * 
the  precipitation  is  complete  on  boiling.  The  compound  thrown 
down  under  these  circumstances  consists  mainly  of  the  di-hydrated 
oxide,  A12O3)2H2O  ;  thus— 
A12O3 ,  Na2O  +  6NH4C1  =  6NaCl  +  6NH3  +  H2O  +  A12O3 ,  2H2O 

BaCO,  suspended  in  water,  precipitates  A12(HO)6,  carbon  di- 
oxide being  evolved.  The  precipitation  is  complete  even  in  the 
cold.f  If  alum  or  aluminium  sulphate  is  used,  the  precipitate  is 
mixed  with  insoluble  barium  sulphate — 

A12(S04)3  +  3BaC03  +  3H2O  =  A12(HO)6  +  3BaSO4  +  3CO2  J 

K  CO  and  Na,CO;;  precipitate  an  uncertain  mixture  of  the 
hydroxide  and  basic  carbonates. 

(NH|)  ,S  precipitates  aluminium  hydroxide,  with  evolution  of 
sulphuretted  hydrogen  (compare  Fe). — 
A12(S04)3  +  3(NH4)2S  +  6H20  -  A12(HO)G  +  3(NH4)2SO4  +  3H2S  \ 

[Aluminium  forms  no  sulphide  in  the  wet  way.     A12S3  (obtained 
by  the   union  of  Al  and  S)  is  decomposed    instantly   by  water, 
forming  the  trioxide,  and  evolving  H2S.] 
Chromium,  Cr. 

DRY  REACTIONS. — Chromium  compounds  impart  to  a  borax 
bead  a  grass-green  colour,  when  heated  either  in  the  outer  or  inner 
blowpipe  flame. 

*  The  determining  cause  of  this  action  of  ammonium  chloride  is  doubtless 
the  instability  of  ammonium  aluminate,  and  the  readiness  with  which  it  under- 
goes hydrolysis.  The  ionised  sodium  aluminate  and  ammonium  chloride  may 
be  regarded  as  first  undergoing  "double  decomposition"  forming  ammonium 
aluminate  which  is  immediately  hydrolysed,  thus — 

6Na  ,  2A1O3  +  6NH4 ,  6C1  =  6Na ,  6cT+  6NH4 ,  2A1O3 
6NH4 ,  2A1O3  =  6NH3  +  H2O  +  H4A12O5  (or  A12O3 ,  2H2O) 

f  In  the  presence  of  certain  organic  acids,  as  oxalic,  tartaric,  or  citric  acids, 
aluminium  hydroxide  is  only  more  or  less  imperfectly  precipitated  by  the 
above-mentioned  reagents,  owing  to  the  formation  of  soluble  double  salts 
of  the  organic  acid  with  aluminium  and  the  alkali  metal ;  such,  for  example, 
as  the  double  tartrate  of  aluminium  and  sodium,  Na2(C4H4O6),Al2(C4H4O6)3. 
This  applies  also  in  the  case  of  the  corresponding  chromium  and  iron  com- 
pounds. 

J  In  these  equations  simple  aluminium  sulphate  is  given  instead  of  alum, 
in  order  not  to  unnecessarily  load  the  equation  with  materials  taking  no  part 
in  the  reaction. 


4O  Qualitative  Analysis. 

When  fused  in  a  platinum  capsule  with  five  or  six  times  their 
weight  of  a  mixture  consisting  of  i  part  of  KNO3  and  2  parts  of 
dry  Na2CO3  or  K2CO3  (or  i  part  of  KC1O3  with  6  parts  of  Na2CO3), 
chromium  compounds  are  converted  into  alkaline  chromates,  which 
appear  as  a  yellow  mass,  soluble  in  water  to  a  yellow  solution. 
In  the  case  of  chromic  oxide,  for  instance,  Cr2O3,  the  reaction  is  the 
following: — 

Cr2O3  +  2K2CO3  +  KC1O3  =  2K2CrO4  +  KC1  +  2CO2 

The  chief  natural  source  of  chromium  is  the  mineral  chrome  iron 
ore,  Cr2O3,FeO.  When  this  is  fused  with  either  of  the  above 
mixtures,  the  same  reaction  takes  place  as  regards  the  chromium, 
while  the  iron  is  changed  to  Fe2O3 ;  thus — 

6Cr2O3,  FeO  +  I2K2CO3  +  7KC1O3  =  i2K2CrO4  +  3Fe,O3  + 

yKCl  +  »  CO, 

[Sodium  peroxide,  Na2O2,  may  be  substituted  as  the  oxidizing 
material,  in  which  case  the  fusion  should  he  carried  out  in  a  silver 
capsule.] 

WET  REACTIONS.— The  two  best-known  classes  of  chromium 
salts  are  derived  from  the  two  oxides,  namely — 

Chromium  sesquioxide  (or  chromic  oxide),  Cr2O3 
Chromium  trioxide  (chromic  anhydride),  CrO3 

Chromic  oxide,  Cr2O3,  is  basic,  uniting  with  acids  to  form  the 
chromic  salts^  such  as  chromic  hydroxide,  Cr2(HO)6  or  Cr2O3,3H2O  ; 
chromic  chloride,  CrCl3 ;  chromic  sulphate,  Cr2(SO4)3 ;  double  potas- 
sium and  chromium  sulphate  (chrome  alum\  K2SO4,Cr2(SO4)3, 
24H2O.  Of  these  salts,  the  hydroxide  alone  is  insoluble  in  water. 

Chromium  trioxide,  CrO3,  is  the  anhydride  of  the  hypo- 
thetical chromic  acid,*  H2CrO4,  which  gives  rise  to  salts  known  as 
chromateS)  analogous  in  constitution  to  the  sulphates. 

Chromates  of  the  metals  of  Groups  IV.  and  V.  are  all  soluble  in 
water,  except  BaCrO4.  The  other  chromates  are  insoluble. 

o.  Chromic  Salts.— These  salts  are  mostly  of  a  purplish 
or  violet-grey  colour  when  solid,  giving  either  a  purple  or  green 
solution  when  dissolved,  the  colour  depending  upon  the  conditions 
of  solution.  Thus  chrome  alum  dissolved  in  cold  water  gives  a 

*  At  first  it  may  confuse  students  to  find  that  chromic  anhydride,  and 
chromic  acid  with  its  salts,  should  not  be  in  the  class  of  chromic  compounds. 
It  must  be  remembered  that  the  classification  is  not  based  upon  the  nomen- 
clature of  the  substances.  "Chromic"  compounds  are  those  containing 
chromium  as  the  "  base,"  or  the  positive  radical  ;  while  in  those  compounds 
derived  from  CrO3  the  element  is  in  the  "acidic"  or  negative  group.  They 
may,  therefore,  be  conveniently  distinguished  as  "  chromic  acid  "  compounds. 


Group  III. — Division  A.  41 

purple  solution,  which  on  boiling  turns  green,*  and  on  long  standing 
again  becomes  purple. 

1711,110  produces  a  bluish  or  greenish-grey  precipitate  of 
chromic  hydroxide,  Cr2(HO)6t,  partially  dissolved  by  excess  of 
ammonia  in  the  cold,  giving  a  lilac-coloured  liquid,  but  completely 
precipitated  on  gently  boiling.  Cr2(HO)(i  is  readily  soluble  in  acids.J 

KHO  and  NaHO  precipitate  Cr2(HO)6,  readily  soluble  in 
excess,  giving  a  deep  green  solution.§  Reprecipitated  by  neutraliza- 
tion with  HC1,  and  by  boiling  with  NH4C1,  as  in  the  case  of  Al. 

BaCO3  precipitates  a  mixture  of  the  hydroxide  and  basic 
carbonate.  Complete  precipitation  only  after  some  hours. 

K.CO5  and  Na2CO3  give  a  similar  precipitate,  the  composition 
of  which  varies  with  the  conditions  of  precipitation. 

(NH4),S  precipitates  Cr2(HO)6.  Precipitation  complete.  [Cr, 
like  Al,  is  incapable  of  forming  a  sulphide  in  the  wet  way.] 

Oxidation  of  Chromic  Compounds. — By  means  of  suitable 
oxidising  agents,  chromic  compounds  are  readily  converted  into 
compounds  of  chromic  acid,  the  mechanism  of  the  change  in  all 
cases  being  the  oxidation  of  the  sesquioxide  into  the  trioxide  ;  thus — 

Cr2O3  +  3O  =  2CrO3 

One  method,  namely,  by  fusion  with  oxidising  agents,  has  been 
explained  under  Dry  reactions.  The  Cr2O3  in  that  instance  is 
oxidised  into  the  potassium  salt  of  chromic  acid.  The  oxidation 
may  be  accomplished  in  the  Avet  way  by  the  following  reactions  : — 

(1)  Boiling  chromic  hydroxide  with  potassium  hydroxide  and 
manganese  dioxide — 

Cr2(HO)6  +  4KHO  +  6MnO2  =  2K2CrO4  +  3Mn2O3  +  5H2O 

(2)  By  substituting  lead  peroxide  for  manganese  dioxide — 
Cr2(HO)6  +  4KHO  +  3PbO2  =  2K2CrO4  +  sPbO  +  5H2O 

In  this  case  secondary  reactions  take  place,  for  PbO  is  soluble 
in  KHO,  and  the  solution  so  formed  then  reacts  upon  the  K2CrO4, 
producing  PbCrO4,  which  also  dissolves  in  KHO. 

*  The  green  colour  is  said  to  be  due  to  the  formation  of  a  basic  salt,  by  the 
action  of  water  upon  the  normal  compound. 

f  The  composition  of  the  precipitate  depends  on  the  conductions  under 
which  it  is  formed.  There  are  several  hydrated  chromic  oxides  ;  compare 
also  Al. 

J  For  the  influence  exerted  by  the  presence  of  organic  acids,  see  footnote, 

P-  39- 

§  The  soluble  compounds  produced  are  similar  to  those  given  by 
aluminium  ;  the  potassium  salt  has  the  composition  Cr2O3,K2O  or  K2Cr2O4. 
They  are  known  as  chromites.  Chrome  iron  ore  is  ferrous  chromite,  Cr2O3,  FeO. 
Although  analogous  to  the  alluminato,  they  must  not  be  called  chromates,  as 
this  name  is  reserved  for  the  salts  of  chromic  acid. 


42  Qualitative  Analysis. 

(3)  By  the  action   of  hypochlorites  (or  hypobromites)    in  the 
presence  of  caustic  alkalies,  either  employed  as  such,  or  formed  in 
the  solution  by  the  use  of  chlorine  or  bromine  in  the  presence  of 
the  caustic  alkali — 

Cr2(HO)6  +  4KHO  +  3KC1O  =  3KC1  +  2K2CrO4  +  5H2O 

(4)  By  the  action  of  sodium  peroxide.     If  a  small  quantity  of 
Na2O2  be  added  to  chromium  hydroxide  suspended  in  water,  and 
the  mixture  gently  warmed,  the  chromium  compound  is  immediately 
converted  into  the  yellow  sodium  chromate  ;  thus — • 

Cr2(HO)6  +  3Na2O2  =  2Na2CrO4  +  2NaHO  +  2H2O 

£.  Chromic  Acid  and  Chromates. — The  acid,  H2CrOt,  has 
never  been  isolated.  The  anhydride,  CrO3,  is  readily  obtained  by 
adding  strong  H2SO4  to  a  cold  strong  solution  of  potassium  di- 
chromate,  when  the  oxide  is  deposited  in  the  form  of  red  silky  needles. 
It  forms  two  classes  of  salts,  viz.  the  normal  chromates,  of  which 
K2CrO4  is  a  type  ;  and  the  dichromates,*  of  which  K2Cr2O7  is  a 
familiar  example.  The  chromates  are  mostly  yellow  or  red  in 
colour,  and  those  which  are  soluble  in  water  (see  p.  40)  impart  a 
yellow  or  orange  colour  to  the  liquid.  The  most  important  of  the 
insoluble  chromates  made  use  of  in  analysis,  and  which  are  all 
precipitated  by  the  addition  of  potassium  chromate  to  solutions  of 
the  metallic  salts,  are  the  following  : — 

Barium  chromate,  BaCrO4  (see  Ba  reactions,  p.  31). 

lead  chromate,  PbCrO4  (see  Pb  reactions,  p.  81).— PbCrO4 
melts  without  decomposition,  and  solidifies  on  cooling  to  a  brown 
crystalline  mass.  At  higher  temperatures  it  gives  off  oxygen — 

2PbCrO4  =  Cr2O3  +  2PbO  +  30 

PbCrO4  (known  as  chrome  yellow),  when  digested  with  NaHO,  or 
with  K2CrO4,  is  converted  into  a  red  basic  lead  chromate  (known 
as  chrome  red] — 

2PbCrO4+  2NaHO  =  Na2CrO4+  H2O  +  Pb2CrO5  (or  PbCrO4,PbO) 
2PbCr04  +  K2CrO4  =  PbCrO^Pb'O  +  K2CrO4,CrO3 

*  The  constitution  of  the  dichromates  (sometimes  wrongly  called  ^/chro- 
mates) may  be  expressed  thus,  K2CrO4,CrO3.  They  are  strictly  analogous 
to  the  pyrosulphates,  K2S2O7,  or  K^SO^SOs,  and  on  this  account  should 
consistently  be  named  pyrochromates.  By  the  action  of  strong  acids,  the 
normal  potassium  chromate  is  converted  into  the  dichromate  ;  thus,  2K2CrO4 
+  H2SO4=  K2SO4+  H2O  +  K2CrO4,CrO3.  And  the  dichromate  is  re-con- 
verted into  the  normal  salt  by  the  action  of  potash — 

K2Cr04,Cr03  +  2KHO  =  2K2CrO4  +  H2O 


Group  III. — Division  A.  43 

Silver  chromate,  Ag2CrO4.  — A  dark  chocolate-red  precipitate, 
soluble  in  ammonia  and  nitric  acid. 

Mercurous  chromate  (basic),  Hg2CrO4,Hg2O.— A  brick- 
red  precipitate,  which,  when  dried,  and  heated  in  a  tube,  gives  a 
mercury  sublimate,  evolves  oxygen,  and  leaves  a  residue  of  Cr2O3. 

Oxidation  of  Chromic  Acid. — Although  CrO3  is  such  a 
highly  oxygenated  compound,  it  appears  to  be  capable  of  still 
further  oxidation  by  hydrogen  peroxide,  giving  rise  to  a  compound 
which  is  believed  by  some  to  be  perchromic  acid,  HCrO4,  or 
2CrO3,H2O2,  and  by  others  to  be  a  compound  of  CrO3  and  H2O.2 
in  undetermined  proportions.  The  interest  of  the  compound  lies 
in  the  fact  that  it  has  an  intense  azure-blue  colour,  and  its  formation 
affords  an  extremely  delicate  test  for  either  chromic  acid  or 
hydrogen  peroxide.  A  few  drops  of  H2O2  (or  a  few  particles  of 
Na2O2)  are  added  to  half  a  test-tube  of  water,  and  the  mixture 
acidified  with  one  or  two  drops  of  HC1.  A  single  drop  of  potassium 
dichromate  solution  added  to  this  produces  an  intense  blue  colour. 
[The  compound  is  very  unstable  in  aqueous  solution,  but  less  so  in 
ether  ;  therefore,  in  testing  for  very  minute  quantities,  ether  should 
be  added  before  the  dichromate  ;  and  on  shaking  the  mixture, 
the  ethereal  layer  which  rises  to  the  surface  will  be  coloured  blue.] 

Reduction  of  Chromic  Acid. — CrO3  is  a  powerful  oxidising 
agent,  giving  up  oxygen  to  oxidisable  substances,  and  being  itself 
reduced  to  Cr2O3  ;  that  is,  to  the  condition  of  a  "  chromic  "  com- 
pound. Thus,  by  sulphur  dioxide  it  is  reduced  to  chromium 
sulphate — 

2Cr03  +  3S02  =  Cr2(S04)3 

The  same  action  takes  place  in  an  acidified  solution  of  potassium 
dichromate — 

K,Cr207  +  H2S04  +  3H2S03  =  Cr2(SO4)3  +  K2SO4  +  4H2O 

Similarly,  chromic  acid  and  chromates  are.  reduced  by  HC1,  oxidis- 
ing the  hydrogen  of  the  acid,  and  liberating  chlorine,  after  the 
manner  of  peroxides  ;  thus — 

CrO8  +  6HC1  =  CrCl3  +  3H2O  +  30 
K2Cr2O7  +  I4HC1  =  2CrCl3  +  2KC1  +  ?H2O  +  3C12 

On  account  of  this  reaction,  a  mixture  of  potassium  dichromate 
and  hydrochloric  acid  is  capable  of  "  oxidising  "  FeCl2  into  FeCl3  ; 
SnCl2  into  SnCl4  ;  As2O3  into  As2O5.  In  all  cases  of  oxidation  by 
chromic  acid,  the  reduction  of  the  chromic  acid  compound  to  the 
state  of  a  "  chromic "  compound  is  evidenced  by  the  change  of 


44  Qualitative  Analysis. 

colour  from  the  yellow  or  orange  of  the  former,  to  the  green  colour 
of  the  latter.  This  reduction  and  change  of  colour  is  at  once  seen 
by  passing  sulphuretted  hydrogen  through  acidified  potassium 
dichromate — 

K2Cr2O7  +  3H2S  +  8HC1  =  2CrCl3  +  2KC1  +  7H2O  +  38 

Many  organic  substances  also  reduce  chromic  acid,  such  as 
oxalic  acid,  and  alcohol.  Thus,  one  molecule  of  oxalic  acid,  C2H2O4, 
requires  one  atom  of  O  to  convert  it  into  CO2  and  H2O — 

C2H204  +  O  -  2C02  +  H,0 

Potassium  dichromate,  in  being  reduced,  has  three  available  atoms 
of  oxygen  to  give  up  ;  thus,  K2Cr2O7  =  Cr2O3,K2O,3O  (in  the 
presence  of  dilute  acids  the  Cr203  and,  K20  form  salts).  Therefore 
one  molecule  of  K2Cr2O7  can  oxidise  three  molecules  of  oxalic  acid, 
resulting  in  the  evolution  of  six  molecules  of  CO2 ;  thus — 

K2Cr207  +  4H2S04  +  3C2H2O4  =  K2SO4  +  Cr2(SO4)3  +  7H2O 

+  6C02 

If  alcohol  be  added  to  a  mixture  of  potassium  dichromate  and 
sulphuric  acid,  the  alcohol  (C2H6O)  is  oxidised  first  to  aldehyde, 
C2H4O,  and  then  to  acetic  acid  (C2H4O2),  and  the  colour  of  the 
mixture  changes  from  orange-red  to  green. 

Iron,  Fe. 

DRY  REACTIONS. — Iron  compounds  impart  to  a  borax  bead 
heated  in  the  outer  flame,  a  colour  which  appears  chocolate  when 
hot,  and  yellow  when  cold.  After  heating  in  the  reducing  flame, 
the  colour  changes  to  a  bottle-green  (the  green  colour  of  common 
bottle  glass  is  caused  by  the  presence  of  iron).  When  heated  on 
charcoal  with  Na2CO3  in  the  inner  blowpipe  flame,  iron  com- 
pounds become  reduced,  and  a  dark  grey  magnetic  mass  is  obtained. 
If  this  be  washed  with  water  in  a  small  mortar,  and  the  end  of 
a  magnet  applied,  it  will  be  attracted  after  the  manner  of  iron 
filings. 

WET  REACTIONS. — The  salts  of  iron  are  derived  from  the  two 
oxides  FeO  and  Fe2O3.*  They  are  both  basic  oxides,  and  give 
rise  to  two  classes  of  salts,  namely,  ferrous  and  ferric  respectively. 
Ferrous  salts  readily  take  up  oxygen,  and  become  converted  into 

*  The  oxide  known  as  magnetic  oxide  of  iron,  or  ferroso-ferric  oxide, 
Fe3O4  or  Fe2O3,FeO,  yields  a  mixture  of  ferric  and  ferrous  salts. 


Group  III.  —  Division  A.  45 

ferric  compounds  ;  while  the  latter,  under  the  influence  of  suitable 
reducing  agents,  easily  pass  back  again  to  the  ferrous  condition. 

(a)  Ferric  Compounds.  —  The  common  ferric  salts  that  are 
soluble  in  water  are  the  chloride,  FeCl3  ;  nitrate,  Fe2(NO3)6,  and 
sulphate,  Fe.,(SO4).5.  These  all  give  yellowish-brown  solutions. 

NH,HO,  KHO,  and  NaHO  throw  down  a  brown  voluminous 
precipitate  of  ferric  hydroxide,*  Fe2(HO)c,  insoluble  in  excess, 
or  in  NH4Cl.t 

K  CO,,  Na2CO3,  and  BaCO3  give  the  same  precipitate,  CO2 
being  liberated  — 

2FeCl3  +  3Na2CO3  +  3H2O  =  Fe2(HO)6  +  6NaCl  +  sCO2 
The  precipitate  is  soluble  in  a  concentrated  solution  of  K2CO3, 
giving  a  deep  reddish  solution  of  unknown  composition.     On  the 
addition  of  water  the  hydroxide  is  reprecipitated.     (With  BaCO3 
basic  carbonates  are  also  precipitated.) 

(1^114)28  produces  a  black  precipitate  of  ferrous  sulphide. 
The  action  may  be  considered  as  taking  place  in  two  stages  : 
(i)  the  reduction  of  the  iron  to  the  ferrous  state,  and  (2)  the 
formation  of  the  ferrous  sulphide  ;  thus,  using  a  dissected  formula 
for  ammonium  sulphide  — 

(i)    :NHSNH,J  H2S  +  2FeCl3  =  2FeCl2  +  2HC1  +  S 
(2)  (NH4)2S  +  FeCl2  =  FeS  +  2NH4C1 

In  the  first  equation  the  hydrochloric  acid  formed  unites  with 
the  ammonia,  producing  2NH4C1. 

Sulphuretted  hydrogen,  H2S,  brings  about  the  first  stage 
in  the  above  action,  reducing  the  iron  from  the  ferric  to  \hzferrous 
state  with  precipitation  of  sulphur,  but  in  the  presence  of  the  free 
acid  which  is  developed  by  the  action,  ferrous  sulphide  cannot  be 
formed.  [Ferric  sulphide  cannot  be  produced  in  the  wet  way.] 

Potassium  ferrocyanide,  K4Fe(CN)6,  or  K4FeCy6,t  pro- 
duces with  ferric  salts  a  dark  blue  precipitate  {Prussian  blue)  — 

3K4(FeCy6)  +  4FeCl3  -  I2KC1  +  Fe4(FeCy6)3 

*  Several  hydrated  ferric  oxides  are  known,  e.g.  Fe2O3l3H2O  ;  Fe2O3,2H2O  ; 
Fe2O3,H2O.  The  composition  of  the  precipitate  produced  by  alkalies  depends 
upon  the  conditions  of  precipitation, 

f  See  footnote  on  p.  39  as  to  the  influence  of  organic  compounds. 

%  "Cy"  is  a  recognised  and  convenient  symbol  for  the  radical  (CN)  ; 
cyanogen.  The  use  of  this  reagent  as  a  test  for  iron  is  unique,  as  being  the 
only  case  in  which  the  reagent  is  itself  a  compound  containing  the  very  metal 
it  is  employed  to  detect.  The  ferrocyanides  and  the  ferricyanides,  however, 
although  compounds  of  iron,  do  not  yield  on  solution  either  ferrous  or  ferric 


ions,    but   the  complex  anion   FefCNj^..    These  give   no  reaction   with   the 

* 


4-6  Qualitative  Analysis. 

This  test  is  extremely  delicate,  but  where  the  amount  of  iron 
is  very  small,  a  blue  or  greenish  coloration  only  is  produced. 
"  Prussian  blue "  is  insoluble  in  hydrochloric  acid,  but  readily 
dissolves  in  oxalic  acid.  It  is  decomposed  by  NaHO  or  KHO, 
with  precipitation  of  ferric  hydroxide — 

Fe4(FeCy6)3  +  I2KHO  =  2Fe2(HO)r>  +  3K4(FeCyG) 

Potassium  ferricyanide,  K3(FeCyG),  gives  no  precipitate  with 
ferric  salts. 

Potassium  thiocyanate,  K(CN)S,  produces  with  ferric 
salts  a  rich  wine-red  coloration,  owing  to  the  formation  of  ferric 
thiocyanate,  Fe(CNS)3,  which  is  soluble  in  water.  The  colour  of 
this  compound  is  very  intense,  hence  the  reaction  may  be  employed 
to  detect  very  small  quantities  of  iron.* 

Reduction  of  Ferric  to  Ferrous  Compounds.— FernV 
compounds  are  readily  reduced  to  the  ferrous  state  ;  they  are 
therefore  oxidising  agents  of  some  importance.  The  action  of 
(NH4)2S  and  of  H2S  has  been  already  mentioned.  Nascent 
hydrogen  reduces  them  in  the  same  way  ;  therefore,  when  metallic 
iron  is  dissolved  in  HC1  or  H2SO4,  the  salts  produced  are  ferrous 
chloride  and  sulphate  respectively.  Nitric  acid,  on  the  other 
hand,  converts  the  iron  into  the  "  ferric  "  state. 

A  ferric  salt  already  in  solution  is  reduced  by  nascent  hydrogen, 
generated  by  introducing  zinc  into  the  acidified  liquid. 

In  passing  from  FeCl3  to  FeCl2,  one  atom  of  chlorine  is 
available  for  oxidising  purposes,  and  is  capable  of  bringing  about 
such  actions  as  the  following— 

The  "oxidation"  of  stannous  chloride,  SnCl2,  to  stannic 
chloride,  SnCl4. 

The  oxidation  of  sulphurous  acid  or  thiosulphuric  acid  into 
sulphuric  acid  ;  thus — 

2FeQ3  +  H2S03  +  H20  =  H2SO4  +  2HC1  +  2FeCl2 
2FeCl3  +  Na2SSO3  +  H2O  =  Na2SO4  +  2HC1  +  2FeCl2  +  S 

reagents  employed  for  detecting  either  ferrous  or  ferric  ions,  and  therefore 
before  the  iron  in  such  compounds  will  give  any  of  the  ordinary  reactions,  its 
union  with  the  cyanogen  radical  must  be  first  destroyed  (see  Cyanides,  p.  163). 
*  This  is  a  "reversible"  reaction,  and  therefore,  when  equilibrium  is 
established,  there  will  be  present  in  the  liquid  both  ferric  chloride  and  potas- 
sium thiocyanate,  thus,  sKCNS  +  FeCl3  ^±  sKCl  +  Fe(CNS)3.  That  this  is 
so  may  be  proved  by  the  following  experiment :  Add  to  a  little  moderately 
dilute  ferric  chloride  a  small  quantity  of  potassium  thiocyanate  ;  then  dilute 
the  liquid  with  water  so  that  the  intensity  of  the  red  colour  is  greatly  reduced, 
and  divide  it  into  two  portions.  To  one  add  more  ferric  chloride,  and  to  the 
other  add  more  potassium  thiocyanate.  In  each  case  the  liquid  becomes  a 
deeper  red  colour. 


Group  III. — Division  A.  47 

(b]  Ferrous  Compounds. — Ferrous  salts  are  usually  pale  green 
when  crystallised,  and  white  when  anhydrous.  Of  the  common  salts 
the  chloride  and  sulphate  are  soluble.  The  latter  readily  forms 
double  salts  with  the  sulphates  of  the  alkalies  (such  as  ferrous  ammo- 
nium sulphate,  FeSo4,(NH4)2SO4,6H2O),  which  are  also  soluble  in 
water,  and  are  less  readily  oxidised  on  exposure  to  the  air  than  ferrous 
sulphate,  which  they  otherwise  closely  resemble  in  appearance.  % 

NH,HO,  KHO,  and  NaHO  produce  a  precipitate  of  ferrous 
hydroxide,  Fe(HO)2,  which  is  at  first  a  dirty  white  colour,  but  which 
rapidly  turns  first  pale  greenish-grey,  then  a  dirty  grey,  and  finally 
brown,  owing  to  its  oxidation  by  atmospheric  oxygen.  The  presence 
of  ammonium  salts  renders  the  precipitation  incomplete.  The 
precipitate  is  not  soluble  in  excess  of  the  reagents ;  boiling  with 
KHO  turns  it  black,  converting  it  into  Fe3O4. 

K  ,CO ,  and  Na2CO3  give  a  white  precipitate  of  ferrous  carbonate, 
FeCO3,  which  on  exposure  to  the  air  quickly  absorbs  oxygen. 

(NH4)2S  precipitates  black  ferrous  sulphide,  FeS.  Readily 
soluble  in  acids,  with  evolution  of  sulphuretted  hydrogen  ;  insoluble 
in  alkalies.  The  precipitate  in  the  moist  state  is  oxidised  on  exposure 
to  the  air  into  ferrous  and  basic  ferric  sulphate. 

K4(PeCy6)  precipitates  potassium  ferrous  ferrocyanide, 
FeK2(FeCy6),  thus— 

K4(FeCye)  +  FeCl2  =  2KC1  +  FeK2(FeCy6) 
When  the  solutions  are  mixed  in  test-tubes  in  the  ordinary  way, 
the  precipitate  has  a  greenish-blue  colour  ;  but  when  the  reaction 
is  made  in  an  atmosphere  free  from  oxygen,  and  the  solutions  are 
previously  boiled  so  as  to  entirely  expel  all  dissolved  oxygen,  the 
precipitate  is  perfectly  white.  It  rapidly  absorbs  oxygen  and 
becomes  blue,  and  is  also  easily  oxidised  to  "  Prussian  "  blue  by 
nitric  acid  or  chlorine  ;  thus — 

4FeK2(FeCy6)  +  2C12  =  Fe4(FeCy6)3  +  K4FeCy6  +  4KC1 

Potassium  ferricyanide,  K3(FeCy6),  gives,  with  ferrous 
salts,  a  precipitate  of  ferrous  ferricyanide,  Fe3(FeCy6)2  (known  as 
TurnbuWs  blue],  which  is  indistinguishable  by  its  appearance  from 
Prussian  blue — 

2K3(FeCy6)  +  3FeCl2  =  Fes(FeCy0)a  +  6KC1 

The  precipitate  is  insoluble  in  hydrochloric  acid,  but  is  decom- 
posed by  caustic  alkalies,  with  the  precipitation  of  ferrous  hydroxide  ; 
thus— 

Fe3(FeCy6)2  +  6KHO  =  2K3(FeCyG)  +  3Fe(HO)2 


4&  Qualitative  Analysis. 

Oxidation  of  Ferrous  to  Ferric  Compounds.— The  ferric 
salts  being  the  more  stable,  the  ferrous  compounds  undergo  oxida- 
tion even  more  readily  than  the  ferric  salts  become  reduced.  Mere 
exposure  to  the  air  in  many  cases  causes  the  change.  In  analysis 
the  oxidation  is  usually  accomplished  either  by  chlorine  (or  bromine) 
or  by  nitric  acid. 

The  chlorine  may  be  employed  in  the  form  of  its  aqueous 
solution  (chlorine  water],  or  more  conveniently  by  generating  the 
gas  in  contact  with  the  ferrous  compound  by  means  of  hydrochloric 
acid  and  potassium  chlorate.  The  solution  of  the  ferrous  salt  is 
acidified  with  concentrated  HC1,  and  heated.  A  few  particles  of 
potassium  chlorate  are  then  dropped  into  the  mixture,  and  the 
heating  continued  for  a  short  time. 

A  mixture  of  HC1  and  KC1O3  evolves  both  chlorine  and  chlorine 
peroxide  ;  thus — 

4KC1O3  +  I2HC1  =  6H2O  +  4KC1  +  3C1O2  +  gCl 

Both  the  free  chlorine  and  the  chlorine  of  the  chlorine  peroxide 
are  available  for  oxidising  the  ferrous  compound  ;  hence  the  equation 
may  be  simplified  as  follows  : — 

KC1O3  +  6HC1  +  6FeCl2  =  6FeCl3  +  KC1  +  3H2O 

When  the  oxidation  is  accomplished  with  nitric  acid,  the  strong 
acid  is  added,  a  few  drops  at  a  time,  to  the  hot  acidulated  solution 
of  the  ferrous  salt.  The  solution  becomes  dark  in  colour,  and  nitric 
oxide  is  disengaged  ;  thus — 

6FeS04  +  3H2SO4  +  2HNO3  =  3Fe2(SO4)3  +  4H2O  +  2ND 
3FeCl2  +  3HC1  +  HNO3  =  3FeCl3  4-  2H.2O  +  NO 

Unless  the  solution  of  the  ferrous  salt  is  acidified,  a  portion  of 
the  iron  is  converted  into  Fe2O3,  which  is  taken  up,  in  the  case 
of  the  sulphate,  by  the  ferric  sulphate,  forming  insoluble  basic  ferric 
sulphates,  Fe2(SO4)s,  ^Fe2O3. 

SEPARATION  OF  THE  METALS  OF  GROUP  I  HA. 

The  separation  of  the  metals  of  this  subdivision  from  the  other 
metals  of  Group  III.,  and  also  from  those  of  Groups  IV.  and  V.,  is 
based  upon  the  fact  that  their  hydrated  sesquioxides  are  precipi- 
tated by  ammonia  in  the  presence  of  ammonium  chloride.* 

*  The  separation  of  Group  IIlA.  from  Group  IIlR.  by  means  of  NH4HO  is 
not  sharp  and  complete  in  all  cases  (see  Manganese  reactions). 


Separation  of  the   Metals  of  Group  III  A  49 

The   separation  of  the  three  metals  of  this   group  from  each 
other  is  based  upon — 

1.  The  oxidation  of  chromic  oxide  to  chromic  acid  ;  and 

2.  The  solubility  of  aluminium  hydroxide  in  caustic  alkalies. 
To  the  solution  add  NH4C1  in  considerable  quantity  ;  heat  the 

mixture  to  boiling,  and  add  NH4HO  carefully  until  precipitation  is 
complete.  Bring  the  liquid  once  more  "  to  the  boil,"  when,  if  suffi- 
cient ammonia  has  been  added,  the  steam  will  smell  of  it.  Filter 
the  mixture  while  hot.* 


The  precipitate  consists  of  A12(HO)6,  Cr2(HO)6,  and  Fe2(HO)6.  Wash 
the  precipitate,  and  transfer  it  (or  a  portion  of  it)  to  a  test-tube  with 
a  small  quantity  of  water.  Add  to  the  mixture  a  little  sodium  per- 
oxide, and  boil  for  a  moment,  until  the  temporary  effervescence  ceases. 
The  chromium  is  oxidised  to  chromate,  and  the  A12(HO)6  dissolves 
in  the  NaHO,  which  is  formed  by  the  action  of  the  sodium  peroxide 
upon  the  water.  Filter. 


The  filtrate  contains  sodium  chromate,Na2CrO4, 
and  sodium  aluminate,  Al2O3,3Na2O.  The 
former  shows  itself  by  the  yellow  colour. 
Divide  into  two  portions — 

(1)  Acidify    with    acetic    acid,    and    confirm 
chromium  by  special  reactions,  e.g.    lead 
acetate. 

(2)  Acidify   with  dilute  nitric   acid,   and   add 
NH4HO.    A  white  precipitate  of  A12(HO)6 
confirms  aluminium. 


The  residue  consists  of 
Fe2(HO)6.  Dissolve 
in  a  little  hot  dilute 
HC1,  and  confirm 
iron  by  special  re- 
actions, e.g.  K4FeCy6 
or  KCNS.f 


The  following  alternative  methods  of  separation  may  also  be 
used. 

(a)  The  precipitated  hydroxides  are  washed  and  dried.  The 
residue  is  then  mixed  with  at  least  six  times  its  weight  of  fusion 
mixture^  and  fused  in  a  platinum  capsule.  In  this  way  the 
chromium  is  converted  into  alkaline  chromate  ;  a  variable  pro- 
portion of  the  aluminium  into  aluminates. 

*  In  the  regular  course  of  a  complete  analysis,  the  filtrate  obtained  here 
will  contain  the  metals  of  Group  Ills.,  IV.,  and  V. 

f  At  this  stage  in  the  process,  the  iron  will  be  in  the  "ferric"  condition. 
To  ascertain  whether  it  was  originally  present  as  a  "ferrous"  or  "ferric" 
compound,  separate  tests  must  be  made  in  the  solution  before  it  has  been 
subjected  to  the  action  of  either  reducing  or  oxidising  agents. 

J  Fusion  mixture  is  a  mixture  of  Na2CO3  and  K2CO3  in  equivalent  propor- 
tions (or  about  10  parts  Na2CO3  to  13  of  K2CO3).  It  is  used  in  preference  to 
Na2CO3  alone,  because  it  has  the  property  of  melting  more  easily  than  either 
carbonate  separately. 

E 


50  Qualitative  Analysis. 

The  fused  mass  is  then  dissolved  in  water,  and  filtered.  The 
filtrate  is  tested  for  aluminium  and  chromium,  while  the  residue  is 
dissolved  and  tested  for  iron,  as  in  the  foregoing  scheme. 

(&)  The  precipitated  hydroxides  are  dissolved  in  a  little  warm 
dilute  HC1,  and  pure  NaHO*  added  in  quantity  considerably 
more  than  sufficient  to  produce  precipitation.  The  mixture  is  then 
boiled  for  a  few  minutes,  and  filtered. 

The  filtrate  contains  sodium  aluminate,  Al2O3,3Na2O.  Add 
dilute  HC1  until  just  acid,  and  reprecipitate  A12(HO)6  with 
ammonia. 

The  precipitate  contains  Cr2(HO)6  and  Fe2(HO)6.  This  is 
dried,  and  fused  with  fusion  mixture.  The  fused  mass  is  dissolved 
in  water  and  filtered.  The  solution  contains  sodium  chromate, 
while  the  Fe2O3  remains  on  the  filter.  These  are  confirmed  as 
in  the  above  methods. 

*  The  commercial  caustic  soda  usually  employed  in  the  laboratory  always 
contains  more  or  less  sodium  aluminate.  The  student  should  test  a  sample  of 
the  reagent  by  neutralising  it  with  HC1,  and  then  adding  NH4HO.  In  the 
method  of  separation  given  above,  this  difficulty  is  avoided,  as  the  sodium 
peroxide  of  commerce  is  usually  free  from  this  impurity. 


CHAPTER  VI. 
REACTIONS  OF  THE  METALS  OF  GROUP  III. — DIVISION  B. 

Manganese,  Mn. 

DRY  REACTIONS. — Manganese  compounds,  when  heated  in  a 
borax  bead  in  the  oxidising  flame,  impart  to  the  bead  a  violet  or 
lilac  colour.  When  heated  in  the  reducing  flame,  the  bead  again 
becomes  colourless. 

A  more  characteristic  reaction  is  based  upon  the  oxidation  of 
manganese  to  manganic  acid.  When  a  manganese  compound  is 
fused  with  KHO,  or  with  Na2CO3  and  a  little  KNO3  or  KC1O3 
upon  a  platinum  capsule,  the  manganese  undergoes  oxidation,  and 
a  deep  green-coloured  mass  is  obtained,  consisting  of  manganates 
of  the  alkali  metals — 

MnO2  -f  Na2CO3  +  O  (from  KC1O3  or  KNO3)  =  Na2MnO4  +  CO2 
MnO2  +  2KHO  +  O  =  K2MnO4  +  H2O 

The  green  mass  (especially  when  obtained  by  fusion  with  KHO) 
dissolves  in  a  small  quantity  of  cold  water  to  a  deep  green  solution. 
When  this  is  either  acidified,  or  warmed,  or  even  largely  diluted 
with  water,  its  colour  changes  from  green  to  pink,  owing  to  the 
conversion  of  the  manganate  into  permanganate  ;  thus — 

3K2MnO4  +  2H2O  =  2KMnO4  +  4KHO  +  MnO2 

WET  REACTIONS. — Of  the  oxides  of  manganese  two  are  basic  : 
the  monoxide,  MnO,  giving  rise  to  the  manganous  salts  ;  and  the 
sesquioxide,  Mn2O3,  to  the  manganzV  salts. 

[The  oxide,  Mn3O4,  yields  both  manganous  and  manganic  salts  ; 
while  MnO2  gives  manganous  salts  with  the  elimination  of  available 
oxygen.] 

The  manganic  salts  are  extremely  unstable  in  solution,  and  are 
incapable  of  existing  as  such  under  the  ordinary  conditions  of 
analysis.  Thus  the  chloride,  believed  to  have  the  composition 
Mn2Cl6,  passes  at  ordinary  temperatures  into  MnCl2  and  chlorine. 


52  Qualitative  Analysis. 

Of  the  common  manganous  salts,  the  chloride,  MnCl2,  4H2O, 
and  the  sulphate,  MnSO4,5H2O,  are  soluble  in  water.  In  the 
crystallised  state  they  have  all  a  pink  colour.  Manganous  salts 
which  are  soluble  in  water  do  not  undergo  atmospheric  oxidation. 

NH,HO,  KHO,  and  NaHO  produce  a  white  precipitate  of 
manganous  hydroxide,  Mn(HO)2.  Insoluble  in  excess  of  the  re- 
agent, the  precipitate  quickly  absorbs  oxygen,  and  is  converted 
into  hydrated  manganic  oxide,  Mn2O3,H2O,  having  a  brown  colour. 

Freshly  precipitated  Mn(HO)2,  while  still  white,  is  soluble  in 
NH4C1,  forming  the  soluble  double  salt  MnCl2,2NH4Cl,H2O  ;  there- 
fore in  the  presence  of  NH4C1,  manganous  hydroxide  is  not  precipi- 
tated by  NH4HO,  and  only  incompletely  by  KHO  or  NaHO. 

The  ammoniacal  solution  of  the  double  chloride,  however,  is 
capable  of  absorbing  oxygen  just  as  the  precipitated  Mn(HO2)  does, 
and  the  liquid  quickly  becomes  muddy,  owing  to  the  precipitation 
from  it  of  the  brown  hydrated  manganic  oxide.* 

X2CO3,  NaaCOs,  or  (NH4)2CO3  give  a  white  precipitate  of 
manganous  carbonate,  MnCO3 ;  insoluble  in  excess.  The  precipi- 
tation is  not  complete  in  the  presence  of  NH4C1.  Manganous 
carbonate  absorbs  atmospheric  oxygen,  and  is  slowly  changed  into 
the  brown  hydrated  oxide. 

(NH4)2S  precipitates  manganous  sulphide,  MnS,  as  a  pale 
pinkish-white  compound,  easily  soluble  in  dilute  acids  (distinction 
from  Ni  and  Co),  soluble  also  in  acetic  acid  (distinction  from  Zn). 
Precipitation  with  (NH4)2S  is  only  complete  in  the  presence  of 
NH4C1.  Owing  to  the  ready  solubility  of  MnS  in  acids,  H2S  is 
incapable  of  precipitating  manganous  sulphide  from  neutral  solu- 
tions ;  for,  by  double  decomposition,  the  acid  of  the  manganous 
salt  would  be  set  free,  and  would  immediately  redissolve  the 
sulphide. 

Manganese  Compounds  as  Oxidising  Agents.— When 
manganese  dioxide  is  acted  upon  by  acids,  a  manganous  salt  is 

*  It  is  this  characteristic  property  of  the  double  manganous  ammonium 
chloride  that  renders  the  complete  separation  of  this  metal  from  those  of 
Group  IIlA.  extremely  difficult  to  accomplish.  The  hydrated  oxides  of  Al, 
Cr,  and  Fe  are  sure  to  carry  down  more  or  less  of  the  manganese  along  with 
them,  and  a  small  quantity  of  manganese,  in  the  presence  of  a  large  proportion 
of  iron,  might  in  this  way  be  altogether  overlooked.  If  the  precipitation  with 
NH4HO  of  Group  III  A.,  be  made  as  quickly  as  possible  in  a  hot  solution, 
and  the  excess  of  ammonia  at  once  boiled  off,  and  the  liquid  filtered  imme- 
diately, the  risk  of  precipitating  the  manganese  maybe  reduced  to  a  minimum. 
The  student  will  do  well  to  practise  the  separation  of  manganese  from  the 
metals  of  Group  IIlA.  by  using  solutions  containing  known  small  propor- 
tions of  a  manganous  salt,  mixed  with  large  quantities  of  iron  or  chromium 
or  aluminium. 


Group  III. — Division  B.  53 

formed,  and  available  oxygen  is  eliminated,  which  either  appears 
as  free  oxygen  gas,  or  as  the  product  of  the  oxidation  of  the  acid  ; 
thus— 

MnO,  +  H2SO4  =  MnSO4  +  H2O  +  O 
Mn62  +  4HC1  -  MnCl2  +  2H2O  +  C12 

The  other  oxides  of  manganese,  higher  than  the  monoxide,  may 
be  regarded  as  compounds  of  MnO2  with  MnO  in  different  propor- 
tions, and  when  acted  upon  by  hydrochloric  acid  they  also  yield 
manganous  chloride  and  chlorine,  the  amount  of  chlorine  being  the 
measure  of  the  amount  of  MnO2  in  the  compound — 

Mn3O4  or  2MnO,MnO2  +  8HC1  =  3MnCL>  +  4H2O  +  C12 
Mn2O3  or  MnO,MnO2  +  6HC1  =  2MnCl2  +  3H2O  +  C12 

Nitric  acid  (free  from  nitrous  add']  has  no  action  upon  man- 
ganese dioxide  ;  when,  therefore,  one  of  these  other  oxides  is  acted 
upon  by  nitric  acid,  the  monoxide  is  converted  into  the  nitrate,  and 
the  dioxide  is  left ;  thus — 

Mn2O3  or  MnO,MnO2  +  2HNO3  =  Mn(NO3)2  +  H2O  +  MnO2 

The  manganates  and  permanganates  are  still  more  powerful 
oxidising  agents.  When  these  are  acted  upon  by  sulphuric  acid,  we 
may  suppose  that  the  hypothetical  manganic  and  permanganic 
acids  are  first  liberated,  which,  being  incapable  of  existence,  break 
up  into  the  unstable  oxides  MnO3  and  Mn2O7 ;  and  that  these  in 
contact  with  sulphuric  acid  form  manganous  sulphate,  with  the 
evolution  of  oxygen  ;  thus— 

fK2MnO4  +  H2SO4  =  K2SO4  +  H2MnO4(=  H2O  +  MnO3) 
I      MnO3  +  H2SO4  =  MnSO4  +  H2O  +  O2 
2KMnO4*+  H2SO4  =  K2SO4  +  2HMnO4(  =  H2O  +  Mn2O7) 
Mn2O7  +  2H2SO4  =  2MnSO4  +  2H2O  +  50 

From  these  equations,  it  will  be  seen  that  from  one  molecule  of 
the  manganate  two  atoms  of  oxygen  are  given  off,  while  two  mole- 
cules of  the  permanganate  evolve  five  atoms  of  oxygen.  When 
acted  upon  by  hydrochloric  acid,  the  equivalent  quantity  of  chlorine 
is  evolved — 

KMnO4  +  8HC1  =  KCL+  MnCl2  +  4H2O  +  SCI 

Potassium  permanganate  is  therefore  a  most  important  oxidising 
material.  In  the  presence  of  hydrochloric  or  sulphuric  acid,  it  is 

*  Some  chemists  prefer  to  use  the  double  formula,  KoMrigOg,  for  potassium 
permanganate. 


54  Qualitative  Analysis. 

capable   of  oxidising   almost    every   oxidisable   compound.     The 
following  may  be  taken  as  typical  examples  : — 

(1)  Sulphurous  acid  to  sulphuric  acid — 

2KMnO4  +  5H2SO3  =  K2SO4  +  2MnSO4  +  3H2O  +  2H2SO4 

(2)  Oxalic  acid  to  carbon  dioxide  and  water— 

H2C2O4  +  O  =  H2O  +  2CO2 

One  atom  of  oxygen  is  required  for  the  oxidation  ot  one  mole- 
cule of  oxalic  acid  ;  therefore  the  five  atoms  of  oxygen  derivable 
from  two  molecules  of  permanganate  will  oxidise  five  molecules  of 
oxalic  acid. 

(3)  Ferrous  sulphate  to  ferric  sulphate,  in  the  presence  of  free 
sulphuric  acid — 

2FeSO4  +  H2SO4  +  O  =  H2O  +  Fe2(SO4)3 
One  molecule  of  KMnO4  is  therefore  capable  of  oxidising  five 
molecules  of  FeSo4  ;  or,  in  the  presence  of  HC1,  of  oxidising  five 
molecules  of  FeCl2  into  FeCl3.* 

In  all  these  cases  of  oxidation  the  change  is  accompanied  by 
the  destruction  of  the  violet  colour  of  the  permanganate  ;  hence  it 
is  perfectly  easy  to  watch  the  progress  of  the  action,  and  to  ascer- 
tain the  exact  moment  when  the  process  is  complete. 

Zinc,  Zn. 

DRY  REACTIONS.— Zinc  compounds  give  no  characteristic 
borax  bead. 

When  heated  on  charcoal  with  sodium  carbonate  in  the  reducing 
flame,  zinc  compounds  are  reduced,  but  the  metal  is  too  volatile  to 
appear  in  the  form  of  globules.  As  it  is  reduced,  it  volatilises  ;  and 

*  Two  powerful  oxidising  substances  are  often  able  to  oxidise  each  other  in 
such  a  way  that  at  first  sight  it  would  seem  that  they  were  playing  the  part  of 
reducing  agents.  Thus,  when  potassium  permanganate  and  hydrogen  peroxide 
are  brought  together,  both  compounds  are  reduced.  The  available  oxygen 
from  the  permanganate  oxidises  the  available  oxygen  of  the  hydrogen  peroxide, 
with  the  result  that  a  number  of  oxidised  atoms  of  oxygen  (i.e.  complete  mole- 
cules) are  set  free.  .In  the  presence  of  H2SO4,  the  five  oxygen  atoms  available 
in  two  molecules  of  permanganate  oxidise  five  available  atoms  of  oxygen  con- 
tained in  as  many  molecules  of  hydrogen  peroxide  ;  thus— 

2KMnO4  +  sH2SO4  +  sH2O2  =  sH2O  +  sH2O  +  K2SO4  +  2MnSO4  +  sO2 


In  the  absence  of  the  free  acid,  hydrated  sesquioxide  of  manganese  is  pre- 
cipitated, which  then  becomes  a  catalytic  agent,  being  alternately  oxidised  and 
reduced  again  — 

2KMn04  +  4H202  =  2KHO  +  Mn2O3,H2O  +  2H2O  +  4O2 


/Mn2O3  +  H2O2  =  H2O  +  2MnO2 
\2MnO2  +  H2O2=  H2O  *  O2  +  Mn2Q3 


Group  III.  —  Division  B.  55 

the  vapour  burns  as  it  passes  through  the  outer  flame,  which  thereby 
becomes  tinged  a  bluish-white  colour.  The  zinc  oxide  which  is 
thus  produced,  deposits  as  an  incrustation  upon  the  charcoal,  which 
is  canary-yellow  while  hot,  becoming  white  on  cooling.  Zinc  oxide 
is  not  volatile,  and  therefore  the  incrustation  does  not  disappear 
when  the  oxidising  flame  is  made  to  play  upon  it.  If  the  zinc  oxide 
be  moistened  with  a  drop  of  cobalt  nitrate,  and  again  heated  in  the 
oxidising  flame,  it  assumes  a  green  colour. 

WET  REACTIONS.  —  Zinc  forms  only  one  series  of  salts,  derived 
from  the  only  oxide,  ZnO. 

Of  the  common  salts  the  chloride,  sulphate,  and  nitrate  are 
soluble  in  water. 

NH,HO,  KHO,  or  NaHO  throws  down  a  white  precipitate  of 
Zn(HO)2,  readily  soluble  in  excess  of  the  reagent,  forming  double 
salts  (sometimes  called  zincates\  such  as  ZnNa2O2.*  Moderately 
strong  solutions  of  these  zincates  may  be  boiled  without  undergoing 
any  change,  but  from  dilute  solutions  Zn(HO)2  is  reprecipitated. 
Zn(HO)2is  soluble  in  NH4C1,  owing  to  its  readiness  to  form  soluble 
double  salts  with  the  alkaline  chlorides,  having  the  general  formula 
ZnCl2,2RCl— 

Zn(HO)2  +  4NH4C1  =  ZnQ2,2NH4Cl  +  2NH4HO 

or  Zn  ,  2HO  +  4NH4  ,  4C1  =  Zn  ,  2C1  ,  2NH4  ,  2C1  +  2NH4HO 

(ionised  double  salt) 

Hence  in  the  presence  of  much  ammonium  chloride,  NH4HO  gives 
no  precipitate,  and  the  precipitation  with  KHO  or  NaHO  is  in- 
complete. 

KoCOj,  Na2CO3,  or  (NH4)2CO3  produces  a  white  precipitate  of 
basic  carbonate,  the  composition  of  which  varies  with  conditions 
of  precipitation,  .rZnCO3,  j/Zn(HO)2,  «H2O.f  The  precipitate  is 
soluble  in  excess  of  (NH4)2CO3.  The  presence  of  much  NH4C1 
partially  or  entirely  prevents  the  precipitation. 

*  Upon  this  property  is  based  the  separation  of  zinc  from  Mn,  Ni,  and  Co, 
but  the  solubility  of  Zn(HO)2  in  caustic  alkalies  is  rendered  less  easy  by  the 
presence  of  the  hydroxides  of  Mn,  Ni  and  Co,  owing  to  the  tendency  of  these 
to  unite  with  the  zinc  oxide  to  form  compounds  which  are  difficultly  decom- 
posed by  the  alkali. 


f  3ZnCl2  +  3Na2CO3  +  sH2O  =  ZnCO3,2Zn(HO)2)H2O  +  6NaCl  +  2CO2 
5ZnCl2  +  5Na2CO3  +  8H2O  =  2ZnCO3,3Zn(HO)2>5H2O  +  loNaCl  +  3CO2 

The  former  of  these  basic  carbonates  is  known  in  pharmacy  as  zinci  carbonas. 
The  normal  zinc  carbonate  is  precipitated  by  hydrogen  sodium  carbonate. 
We  may  suppose  that  the  acid   which  is  in  this  case  liberated,  prevents  the 
formation  of  the  hydroxide  — 

ZnCl2  +  HNaC03  =  ZnCO3  +  HC1  +  NaCl 


56  Qualitative  Analysis. 

The  precipitate  is  soluble  also  in  a  concentrated  solution  of 
K2CO3,  but  is  reprecipitated  on  dilution  with  water  (compare 
Fe,  Ni,  Co). 

(NH4)2S  throws  down  a  white  precipitate  of  zinc  sulphide, 
ZnS.  In  the  presence  of  NH4C1  the  precipitation  is  complete  even 
from  dilute  solutions.  ZnS  is  soluble  in  dilute  mineral  acids,  hence 
H2S  is  incapable  of  completely  precipitating  this  sulphide  from 
neutral  solutions  of  the  zinc  salts  of  such  acids — 

ZnCl2  +  H2S  ^  ZnS  +2HC1 

ZnS  is  insoluble  in  acetic  acid  (contrast  MnS),  therefore  from  the 
acetate,  or  other  zinc  salts  in  presence  of  an  alkaline  acetate,  ZnS 
is  completely  precipitated  by  H2S  ;  thus — 

Zn(C2H302)2  +  H2S  =  ZnS  +  2H(C2H3O2) 
ZnCl2  +  2Na(C2H3O2)  +  H2S  =  ZnS  +  2NaCl  +  2H(C2H3O2) 

Nickel,  Ni. 

DRY  REACTIONS.— Nickel  compounds  impart  a  dark  red-brown 
colour  to  the  borax  bead  when  heated  in  the  oxidising  flame,  the 
colour  becoming  brownish-yellow  on  cooling.  In  the  reducing 
flame  the  borax  bead  becomes  opaque  and  grey.  In  a  bead  of 
microcosmic  salt,  the  red-brown  colour  persists  in  both  flames. 

The  presence  of  other  colour-producing  oxides  renders  this  test 
uncertain,  while  even  traces  of  cobalt  entirely  mask  it.  Heated 
on  charcoal  with  Na2CO3)  metallic  nickel  is  obtained  as  a  grey 
feebly  magnetic  mass. 

WET  REACTIONS.— Only  one  of  the  oxides  of  nickel  is  basic, 
namely  NiO,  hence  only  one  series  of  salts  exists.  The  sesquioxide, 
Ni2O3,  behaves  like  a  peroxide  ;  thus — 

Ni2O3  +  2H2SO4  =  2NiSO4  +  2H2O  +  O 

In  the  crystalline  or  hydrated  condition  the  nickel  salts  have  a 
green  colour,  and  dissolve  to  green  solutions.  The  anhydrous  salts 
are  pale  yellow.  Of  the  common  salts,  the  chloride,  nitrate,  and 
sulphate  are  soluble  in  water.  Nickel  salts  form  a  number  of 
soluble  double  salts  ;  those  with  ammonium  salts  being  important. 
EHO  or  NaHO  gives  a  pale  bluish-green  precipitate  of  nickelous 
hydroxide,  Ni(HO)2)*  insoluble  in  excess  of  either  reagent;  soluble 

*  The  composition  of  the  precipitate  is  more  exactly  expressed  by  the 
formula  4Ni(HO)2,H2O.  When  dried  and  strongly  heated,  it  is  converted 
into  NiO. 


Group  III — Division  B.  57 

in  ammonium  salts.  Ni(HO)2  is  not  oxidised  on  exposure  to  air, 
but  it  is  converted  into  black  hydrated  sesquioxide,  Ni2O3,3H2O, 
by  hypochlorites,  or  by  the  action  of  chlorine  in  the  presence  of 
caustic  alkalies  ;  thus — 

2Ni(HO)2  +  NaCIO  +  H2O  =  NaCl  +  Ni2O3,3H2O 
2Ni(HO)2  +  2NaHO  +  C12  =  2NaCl  +  Ni2O3,3H2O* 

NH  HO  forms  a  number  of  readily  soluble  double  compounds 
with  nickel  salts.  When  added  to  an  acid  solution  of  a  nickel  salt, 
e.g.  NiSO4,  no  precipitate  is  produced,  owing  to  the  ready  solubility 
of  Ni(HO)2  in  ammonium  salts.  With  neutral  solutions  partial 
precipitation  takes  place,  the  precipitate  quickly  dissolving  in  excess 
of  ammonia  to  a  greenish-blue  solution.  This  blue  liquid  contains 
in  solution  the  salt  NiSO4,4NH3.  The  nickel  in  this  solution  is 
not  oxidised  by  hypochlorites,  but  it  is  completely  precipitated  as 
Ni(HO)2  by  KHO. 

K,CO  or  Na2CO3  produces,  a  pale-green  precipitate  of  basic 
carbonate,  ^NiCO3,jNi(HO)2. 

The  precipitate  is  soluble  to  a  pale-green  solution  in  a  concen- 
trated solution  of  K2CO3,  but  is  reprecipitated  on  dilution  with  water. 

(NH4)2CO3  gives  no  precipitate  in  acid  solutions,  but  from 
neutral  solutions  a  similar  compound  is  produced,  which  dissolves 
in  excess  of  the  reagent  to  a  bluish  solution. 

(NH4)2S,  or  H2S  in  presence  of  ammonia,  produces  a  black  pre- 
cipitate of  NiS,  soluble  to  a  slight  extent  in  excess ;  more  readily 
soluble  if  ammonia  or  polysulphides  of  ammonia  are  present, 
yielding  a  brown  solution.  From  this  solution  the  dissolved  NiS 
is  reprecipitated  slowly  on  boiling,  more  quickly  after  acidifying 
with  acetic  acid,  or  the  addition  of  ammonium  acetate. 

NiS  is  only  difficultly  soluble  in  strong  HC1,  and  almost  insoluble 
in  the  dilute  acid  ;  also  in  acetic  acid.  Readily  soluble  in  aqua 
rcg.ia,  or  in  HC1  and  a  crystal  of  KC103,  yielding  NiCl2 ;  soluble 
also  in  HNO3. 

*  This  higher  oxide  undergoes  rapid  alternate  reduction  and  oxidation  in 
the  presence  of  the  hypochlorite.  As  soon  as  it  is  formed  it  is  acted  upon  by 
the  hypochlorite,  with  evolution  of  oxygen  ;  thus,  formulating  the  oxides  for 
simplicity — 

Ni203  +  NaCIO  =  NaCl  +  2NiO  +  O2 

The  compound  therefore  becomes  a  catalytic  agent,  causing  the  evolution  of 
oxygen  from  a  relatively  infinite  quantity  of  a  hypochlorite.  Cobalt  oxide 
behaves  in  the  same  way  (see  Methods  of  obtaining  oxygen,  Newth's  "  Inorganic 
Chemistry  ").  The  hydrated  sesquioxide  behaves  with  acids  in  the  same  way 
as  the  anhydrous  oxide  ;  thus  with  HC1  it  yields  chlorine — 

Ni203,3H20  +  6HC1  =  zNiCl2  +  6H2O  +  C12 


58  Qualitative  Analysis. 

H2S  only  produces  complete  precipitation  of  NiS  from  a  warm 
solution  of  the  acetate,  or  from  other  nickel  salts  in  the  presence  of 
an  alkaline  acetate.  In  the  case  of  neutral  solutions  of  nickel  salts 
with  mineral  acids,  the  precipitation  is  only  partial,  while  in  acid 
solutions  it  does  not  take  place  at  all. 

Cobalt,  Co. 

DRY  REACTIONS. — Cobalt  compounds  impart  to  a  borax  bead 
a  rich  blue  colour,  when  heated  either  in  the  oxidising  or  reducing 
flame.  The  test  is  characteristic  and  delicate.  Many  metallic 
oxides  which  give  coloured  borax  beads  become  either  colourless  or 
lighter  in  colour  when  heated  in  the  reducing  flame,  hence  when 
mixed  with  these,  the  blue  colour  of  the  cobalt  becomes  more 
visible  after  heating  the  bead  in  the  inner  flame. 

WET  REACTIONS. — Cobalt  forms  a  number  of  oxides,  two  only 
of  which  are  basic.  The  cobaltous  salts  are  derived  from  CoO, 
while  the  feebly  basic  sesquioxide  Co2O3  forms  the  very  unstable 
cobaltic  salts. 

Of  the  common  cobaltous  salts,  the  sulphate,  nitrate,  and  haloid 
salts  are  soluble  in  water.  In  the  hydrated  condition  they  have  a 
pink  colour,  dissolving  to  pink  solutions.  In  the  anhydrous  state 
they  are  blue.  A  strong  aqueous  solution  (pink)  will,  however,  turn 
blue  when  boiled,  and  return  to  its  original  pink  colour  when  again 
cooled.  In  alcohol,  cobalt  chloride  dissolves  to  a  deep  blue  solu- 
tion, which  turns  pink  on  the  addition  of  water. 

KHO  or  NaHO  gives  a  greenish-blue  precipitate  of  a  basic  salt. 
On  boiling  with  excess  of  the  alkali,  the  precipitate  is  converted 
into  the  pink  hydroxide,  Co(HO)2,  which,  however,  is  coloured 
more  or  less  brown  by  the  oxidation  of  a  portion  of  it  (by  atmo- 
spheric oxygen)  into  the  hydrated  cobaltic  oxide,  Co2O3,3H2O,  or 
Co2(HO)6. 

Co(HO)2  is  oxidised  by  hypochlorites,  in  the  same  manner  as 
the  corresponding  nickel  compound. 

NH  HO  produces  no  precipitate  in  acid  solutions.  In  neutral 
solutions  it  causes  partial  precipitation  of  a  basic  salt,  which  dis- 
solves easily  in  excess.  The  solution,  which  has  a  brownish  colour, 
absorbs  oxygen,  and  becomes  darker  in  colour.* 

*  Cobalt  (in  common  with  platinum  and  a  few  other  metals)  possesses  the 
property  of  forming  a  large  number  of  compounds  with  ammonia.  These 
salts,  which  are  many  of  them  extremely  complex  in  their  constitution,  and 
difficult  to  classify,  are  known  as  cobaltamines.  The  ammonio-cobaltous  com- 
pounds easily  absorb  oxygen,  and  pass  into  the  more  complex  and  more 
numerous  ammonio-cobaltic  salts.  Nickel  does  not  form  compounds  corre- 
sponding to  the  ammonio-cobaltic  salts. 


Group  III. — Division  B.  59 

K  CO  or  Na2CO3  precipitates  a  lilac-coloured  basic  carbonate, 
.t-CoCO3,  yCo(HO)2. 

The  precipitate  is  soluble  to  a  violet  solution  in  a  concentrated 
solution  of  K2CO3,  but  is  reprecipitated  on  dilution  with  water 
(compare  Co,  Zn,  Cu).  In  all  these  cases  the  solution  is  believed 
to  contain  a  double  carbonate,  which  is  unable  to  exist  in  the 
presence  of  much  water. 

HKCO ,  gives  a  pinkish  precipitate  of  normal  cobalt  carbonate, 
CoCO3.  The  precipitate  so  obtained  is  soluble  in  hydrogen  per- 
oxide, yielding  a  deep  green  solution.  The  production  of  this 
green  solution  constitutes  a  very  delicate  test  for  cobalt.  It  may  be 
carried  out  in  the  following  way  ;  to  a  strong  solution  of  potassium 
(or  sodium)  bicarbonate,  about  an  equal  bulk  of  hydrogen  peroxide 
is  added.  On  the  addition  of  a  single  drop  of  cobalt  chloride,  the 
green  colour  will  appear.* 

(NH4)2CO3  gives  no  precipitate  in  acid  solutions.  In  neutral 
solutions  partial  precipitation  of  basic  carbonate  takes  place  ;  the 
precipitate  readily  dissolves  in  excess  to  a  reddish  solution. 

(NH4)2S  gives  a  black  precipitate  of  cobaltous  sulphide,  CoS. 
The  precipitation  is  complete  in  the  presence  of  NH4C1  (contrast 
NiS).  CoS  is  soluble  in  HNO3,  in  "aqua  regia,"  and  in  HC1  with 
the  addition  of  a  crystal  of  KC1O3 ;  difficultly  soluble  in  strong 
HC1  ;  practically  insoluble  in  dilute  HC1. 

H2S  precipitates  CoS  under  the  same  conditions  as  apply  in  the 
case  of  NiS. 


*  It  has  been  found  (Durrani)  that  salts  of  other  carboxylic  acid  (e.g. 
oxalic,  acetic,  tartaric,  etc.)  yield  similar  green  compounds  when  mixed  with 
hydrogen  peroxide  and  a  cobalt  salt.  It  is  believed  that  all  trfese  compounds 
contain  the  group  >  Co  —  O  —  Co  <  and  that  the  constitution  of  the  compound 
with  hydrogen  potassium  carbonate  is  represented  by  the  formula — 

C02K  •  0^f,        0      rn   ,0  •  C02K 
C02K   •  o>C°-°-Co<0  '  C02K 

Carbon  dioxide  is  evolved  during  its  formation,  as  seen  by  the  equation — 
4HKC03  +  2CoC03  +  H202  =  Co20(KC03)4  +  2CO2  +  H2O 

The  green  colour  of  the  solution,  curiously  enough,  is  indistinguishable  by  the 
eye  from  the  green  of  nickel  salts.  Nickel  is  incapable  of  giving  this  reaction, 
hence  it  constitutes  a  ready  means  of  detecting  even  traces  of  cobalt  in  the 
presence  of  nickel. 


60  Qualitative  Analysis. 

The  Cyanides. 

The  metals  of  Group  IIlB.  all  form  compounds  with  cyanogen, 
some  of  which  are  of  importance  in  analysis.  The  addition  of 
potassium  cyanide  to  solutions  of  either  a  manganous,  zinc,  nickel, 
or  cobaltous  salt  results  in  the  precipitation  of  the  cyanide,  of  the 
general  formula  RCy2.  In  each  case,  also,  the  precipitated  cyanide 
readily  dissolves  in  excess  of  the  KCy,  producing  double  cyanides. 
In  the  case  of  Mn,  Zn,  and  Ni,  these  have  the  general  formula 
2KCy,RCy2,  or  K2RCy4,  while  the  cobalt  double  cyanide  has  the 
formula  4KCy,CoCy2  (or  K4CoCy6,  corresponding  to  potassium 
ferrocyanide,  K4FeCy6). 

From  the  solutions  of  these  double  cyanides,  (NH4)2S  is  incapable 
of  precipitating  the  metals  as  sulphides,  for  the  reason  that  the 
metal  is  present  in  the  complex  anious  RCy4  and  RCy6.  The 
addition  of  dilute  HC1  or  H2SO4  to  these  solutions  causes  the 
reprecipitation  of  the  cyanide— 

2KCy  ,  NiCy2  +  2HC1  =  2KC1  +  2HCy  +  NiCy2  or 

+  +       -          + 

2K  ,  NiCy4  +  2H  ,  2C1  =  2K  ,  2C1  +  2HCy  +  NiCy2 

In  the  case  of  the  Mn,  Zn,  and  Ni  compounds,  boiling  with  dilute 
acid  completely  decomposes  the  metallic  cyanide  ;  thus— 

2KCy,NiCy2  +  4HC1  =  2KC1  +  4HCy  +  NiCl2 

•  The  chief  interest  attaching  to  these  compounds  in  this  connection, 
is  due  to  the  difference  between  the  behaviour  of  cobalt  and  of 
nickel  towards  cyanogen.  The  double  cyanide  of  potassium  and 
cobalt,  4KCy,CoCy2  or  K4CoCy6  (like  the  corresponding  iron- 
compound)  exhibits  a  great  readiness  to  undergo  oxidation,  and 
to  pass  from  the  condition  of  cobaltocyanide  to  cobalticyanide  j 
thus  — 

2K4CoCy6  +  O  +  H2O  =  2KHO  +  2K3CoCy6 

Nickel  forms  no  compounds  corresponding  to  the  cobalticyanides. 
When  the  nickel  double  cyanide  is  oxidised  (e.g.  by  chlorine, 
bromine,  or  hypochlorites),  the  nickel  is  converted  into  the  black 
hydrated  sesquioxide  and  precipitated  ;  thus— 


2K2NiCy4  +  NaCIO  +  5H2O  =  Ni2O3,3H2O  +  NaCl 

+  4HCy 

While   under  the   same    oxidising    treatment,  the   cobalt   double 


Separation  of  the  Metals  of  Group  IllB.          61 

cyanide  is  converted  into  potassium  cobalticyanide,  which  remains 
in  solution,  and  can  therefore  he  separated  by  filtration  ;  thus  — 

2K4CoCy,.  +  NaCIO  +  H2O  =  NaCl  +  2KHO  +  2K3CoCy6 

The  cobaltocyanides  are  so  readily  oxidised,  that  by  merely  boiling 
the  aqueous  solution  the  change  is  effected  ;  thus — 

K4CoCyc  +  H2O  =  KHO  +  K3CoCyc  +  H 

Owing  to  this  reaction,  the  double  cobalt  cyanide  does  not  behave, 
when  boiled  with  dilute  acid,  in  the  same  manner  as  the  nickel 
double  cyanide  (see  above).  We  may  suppose  that  the  cobalt 
cyanide  first  formed,  in  the  presence  of  free  HCy  (liberated  partly 
from  the  double  cyanide  and  partly  from  the  excess  of  KCy 
present)  and  water,  is  converted  into  cobalticyanic  acid  in  the 
following  manner  : — 

(i)  4KCy,CoCy2  +  4HC1  =  4KC1  +  4HCy  +  CoCy2 
(2)  CoCy2  +  4HCy  +  H2O  =  H3CoCyG  +  H 

It  follows  from  this  that  the  addition  of  dilute  HC1  to  a  solution  of 
a  cobalticyanide  does  not  result  in  the  precipitation  of  a  cyanide, 
but  only  in  the  liberation  of  cobalticyanic  acid. 

The  metal  cobalt  in  cobalticyanides,  like  iron  in  ferricyanides, 
is  not  precipitated  by  the  ordinary  reagents  for  these  metals. 

SEPARATION  OF  THE  METALS  OF  GROUP  1 1  IB. 

The  separation  of  the  four  metals  of  this  group  is  based  upon— 

1.  The  solubility  of  Zn(HO)2  in  caustic  alkalies. 

2.  The  solubility  of  MnS  in  acetic  acid. 

3.  The  different  behaviour  of  the  double  cyanides  of  Ni  and  Co, 
when  subjected  to  oxidising  agents. 

In  the  course  of  a  systematic  analysis,  the  metals  of  Group  1 1  IB. 
are  looked  for  in  the  solution  after  Group  I  HA.  has  been  separated 
by  means  of  NH4HO  in  the  presence  of  NH4C1.  H2S  is  passed 
through  this  ammoniacal  solution,  whereby  the  sulphides  of  Mn, 
Zn,  Ni,  and  Co  are  precipitated.  The  mixture  is  gently  warmed 
and  filtered.* 

*  Even  when  the  exercise  is  confined  only  to  Group  IllB.,  the  student  is 
advised  to  go  through  the  step  of  precipitating  the  group  as  sulphides  in  the 
aramoniacal  solution  prepared  as  described  on  p.  48. 


62 


Qualitative  A  nalysis. 


The  precipitate,  consisting  of  MnS,  ZnS,  NiS,  and  CoS,  is  dissolved  in 
hot  dilute  HC1,  with  the  aid  of  a  few  particles  of  KC1O3.*  The  solution 
is  boiled  until  it  no  longer  smells  of  chlorine,  and  NaHO  added  in 
excess.  It  is  again  boiled  (to  ensure  the  solution  of  all  the  Zn(HO)2), 
and  filtered  after  cooling  (hot  NaHO  will  attack  the  filter-paper). 


The     solution 

contains  so- 
dium zincate, 
ZnNa2O2. 


The  precipitate,  consisting  of  Mn(HO)2,  Ni(HO)2, 
and  Co(HO)2,  is  washed  to  remove  the  soluble  zinc 
salt  still  adhering  to  it,  and  then  dissolved  in  the 
smallest  quantity  of  warm  HC1.  The  solution  is 
nearly  neutralised  with  NH4HO,  a  considerable 


Pass       H  2  S 

quantity  of  ammonium  acetate  added,  and  H2S  passed 

through      the 

through  the  mixture  until  precipitation  is  complete. 

liquid  (or  add 

H2S      water), 

• 

when  a  white 

The  solution 

The  precipitate  (NiS  and  CoS).     Test 

precipitate   of 

contains 

a  small    portion  with  a  borax   bead. 

ZnS     is    pro- 
duced. 

manganous 

Blue  colour    indicates    cobalt.      Dis- 

acetate, 

solve    the     precipitate    in    the    least 

Mn(C2H302)2. 

quantity  of  aqua  regia  ;   boil  off  the 



excess  of  acid,  and  nearly  neutralise 

Precipitate 

with  Na2CO3  (avoid  dilution).     Add 

MnCO3  by  add- 

KCy (freshly  made  solution)  until  the 

ing    Na2CO3. 
Filter  ;     wash 

precipitated   cyanides   are  just  redis- 
solved,    add    NaHO   in   considerable 

thoroughly  ; 
dissolve        in 
HC1,  and  con- 

quantity, and  then  bromine  water  f  until 
the   colour   of  the    bromine  persists. 

firm     Mn    by 

Filter. 

precipitation 

of  the  pinkish 

MnS,        with 

The  solution  con- 

The     precipitate 

(NH4)2S  after 

tains          sodium 

consists     of      the 

NH4C1      and 

cobalticyanide, 

black  hydrated  ses- 

NH4HO. 

Na3CoCy6. 

quioxide  of  nickel, 



_..._, 

NLO,  iH,O. 

The    formation 

The  cobalt  may  be 

j.^  ^g^SJ  J       2      ' 

of   the    green 

confirmed   by  the 

Confirm  (i)  by  borax 

m  anganat  e 
may    also    be 
used  as  a  con- 

borax bead,  if  the 
test  was  not  made, 
or  was  not  satis- 

bead ;  (2)  wash  pre- 
cipitate, dissolve  in 
HC1,    and    obtain 

firmatory  test. 

factory,    with   the 

the     characteristic 

sulphides. 

reaction           with 

(NH4)2S. 

*  Should  the  precipitate  not  be  black,  only  MnS  and  ZnS  can  be  present  ; 
and  as  these  are  easily  soluble  in  HC1,  the  KC1O3  should  in  this  case  not  be 
added. 

f  Bromine  is  more  soluble  in  water  containing  KBr  in  solution  than  in  water 
alone.  By  the  use  of  such  a  solution,  therefore,  unnecessary  dilution  is  avoided. 


CHAPTER    VII. 
THE  PHOSPHATES. 

THE  salts  of  phosphoric  acid,  H3PO4,  are  all  of  them  insoluble  in 
water  except  those  of  the  alkali  metals  and  ammonium.  Lithium 
phosphate  is,  however,  only  difficultly  soluble.  The  addition,  there- 
fore, of  a  solution  of  one  of  these  soluble  phosphates  to  a  solution 
of  a  metallic  salt  results  in  the  precipitation  of  the  phosphate  of  the 
metal  contained  in  the  salt ;  thus,  using  hydrogen  disodium 
phosphate  as  being  the  most  convenient  reagent — 

HNa2PO4  +  3AgNO3  =  Ag3PO4  +  2NaNO3  -f  HNO3 

HNa2PO4  +  BaCl2  =  HBaPO4*+  2NaCl 
2HNa2PO4  +  3CaQ2  =  Ca3(PO4)2  +  4NaCl  +  2HC1 

The  most  delicate  reaction  for  the  detection  of  a  phosphate  is 
the  formation  of  a  canary-yellow  crystalline  precipitate  o>i  ammonium 
phosphomolybdate,  upon  the  addition  of  a  nitric  acid  solution  of 
ammonium  molybdate,  (NH4)2MoO4,  to  a  solution  of  a  phosphate.  A 
few  drops  of  the  solution  of  a  phosphate  are  taken,  and  a  very 
large  excess  of  the  nitric  acid  solution  of  ammonium  molybdate  is 
added,  and  the  mixture  slightly  warmed  when  the  yellow  precipitate 
crystallises  out.  It  is  insoluble  in  water  or  acid  provided  a  con- 
siderable quantity  of  ammonium  molybdate  is  present.  The 
precipitate  is  believed  to  have  the  composition  i2MoO3,(NH4)3PO4. 
It  is  decomposed  by  ammonia  into  (NH4)2MoO4  and  (NH4)3PO4. 

When  the  phosphates,  insoluble  in  water,  are  acted  upon  by 
acids,  they  either  dissolve,  or,  if  the  acid  employed  is  capable  of 
forming  an  insoluble  compound  with  the  metal,  they  are  converted 
into  this  compound,  and  phosphoric  acid  is  set  free.  Thus,  calcium 
phosphate  is  decomposed  by  sulphuric  acid,  insoluble  calcium 
sulphate  being  formed,  while  the  liberated  phosphoric  acid  passes 
into  solution — 

Ca3(P04)2  +  3H2S04  =  3CaS04  +  2H3PO4 

*  It  depends  upon  the  conditions  of  precipitation  whether  the  so-called  acid 
phosphate,  HBaPO4,  or  the  normal  salt,  Ba3(PO4)2,  is  produced. 


64  Qualitative  Analysis. 

In  those  cases  where  the  phosphate  is  entirely  dissolved  by  the 
acid,  it  is  reprecipitated  unaltered  by  the  addition  of  ammonia, 
ammonium  sulphide,*  caustic  alkalies,  alkaline  carbonates,  or  in 
some  cases  by  barium  carbonate.  For  example,  the  phosphates 
of  Ni,  Co,  Mn,  Zn,  Ba,  Sr,  Ca,  and  Mg  are  all  soluble  in  HC1  ;  and 
if  ammonia  be  added  to  the  hydrochloric  acid  solution  of  any  one 
of  these,  it  at  once  causes  the  reprecipitation  of  the  phosphate. 
The  significance  of  this  fact,  and  its  bearing  upon  the  method  of 
separation  of  the  metals  of  Group  III.,  will  be  obvious.'  When 
NH4C1  and  NH4HO  are  used  to  separate  Group  IIlA.  from 
Group  1 1  IB.  by  precipitating  the  hydroxides  of  Al,  Cr,  and  Fe, 
they  will  at  the  same  time  throw  down  the  phosphates  of  all  or 
any  of  the  above-named  metals,  if  they  happen  to  be  present  as 
phosphates  in  the  solution  under  examination.  Hence  it  is 
necessary  to  adopt  some  method  for  withdrawing  the  phosphoric 
acid  which  is  in  combination  with  these  metals,  and  of  converting 
them  all  into  such  salts  as  chlorides  before  it  is  possible  to  attempt 
their  separation  on  the  basis  of  the  plans  laid  down  on  pp.  49 
and  62. 

It  will  be  well,  therefore,  at  this  point  to  enter  somewhat  fully 
upon  a  study  of  such  reactions  with  phosphates  as  are  made  use  of 
in  analysis. 

The  Solution  of  Phosphates  by  Acids.  I .  By  Hydrochloric 
Acid. — The  phosphates  of  the  metals  of  Groups  III.,  IV.,  and  V., 
which  are  insoluble  in  water,  are  dissolved  by  hydrochloric  acid. 
When  such  a  phosphate  is  dissolved  in  this  acid,  we  may  consider 
that  double  decomposition  takes  place  between  the  acid  and  the 
phosphate,  just  as  in  the  case  of  the  calcium  phosphate  and 
sulphuric  acid  given  above,  except  that  here  the  metallic  salt 
formed  by  the  union  of  the  metal  with  the  solvent  acid  is  soluble  in 
water.  Both  of  the  products  of  the  interaction  are  ionised,  and 
complete  solution  results  from  the  action.  Thus,  in  the  case  of  the 
solution  of  barium  phosphate  in  hydrochloric  acid  barium  chloride 
and  phosphoric  acid  are  formed — 

HBaP04  +  2HC1  =  BaCl2  +  H3PO4 

The  barium  chloride  and  phosphoric  acid  are  incapable  of  reacting 
upon  each  other,  and  therefore  exist  together  in  the  solution. f 

*  Except  those  of  Fe,  Mn,  Xi,  Co,  Zn.     See  p.  65. 

f  It  is  quite  easy  to  prove,  in  this  particular  case,  that  when  barium  phosphate 
is  dissolved  in  hydrochloric  acid,  the  solution  contains  barium  chloride  and 
phosphoric  acid.  This  may  be  done  by  gently  evaporating  the  solution  down 
upon  a  steam-bath  untfl  crystals  begin  to  separate  out.  If  these  are  drained, 
and  washed  a  few  times  with  strong  hydrochloric  acid  (in  which  BaO2  is  nearly 


The  Phosphates.  65 

Now,  when  an  alkali  is  added  to  such  a  solution  of  a  phosphate 
in  hydrochloric  acid,  as  already  mentioned  the  phosphate  is  at 
once  reprecipitated,  which  may  be  expressed  by  the  equation — 

BaCl2  +  H3PO4  +  2NaHO  =  2NaCl  +  HBaPO4  +  2H2O 

If  expressed  in  the  form  of  an  ionic  equation  it  will  be  seen  that 

+ 
the  equilibrium  of  the  system  is  disturbed  by  the  removal  of  H  and 

IK)  ions  which  unite  to  form  molecules  of  (practically  unionised) 
water — 

Ba,2Cl  +  3H,PO4  +  2Na,2H~O  =  2Na,2Ci  +  HBaPO4  +  2H2O 

In  order,  therefore,  to  formulate  the  various  reactions  which  phos- 
phates in  acid  solutions  undergo,  the  solution  of  the  phosphate 
will  be  expressed  in  the  equations  as  a  mixture  of  phosphoric  acid 
and  the  metallic  salt  of  the  solvent  acid.  In  this  way  the  mechanism 
of  the  reactions  will  be  rendered  easy  of  understanding.*  Two  or 
three  examples  of  the  precipitation  of  phosphates  from  hydrochloric 
acid  solutions  by  various  alkaline  reagents  may  be  given. 

(a}  Precipitation  of  calcium  phosphate  from  HC1  solution  by 
NH4HO,  giving  only  the  final  result — 


:  3CaCl2  +  2H3PO4  :   +  6NH4HO  -  6NH4C1  +  Ca3(PO4)2+  6H2O 

(b]  Precipitation  of  aluminium  phosphate  by  (NH4)2S  — 
:  2Alci3  T2H3PO4  i    +  3(NH4)2S  =  6NH4C1  +  2AlPO4t 


In  the  case  of  the  metals  Fe,  Mn,  Zn,  Ni,  and  Co  (metals  which  are 
capable  of  forming  insoluble  sulphides),  the  phosphates  are  not 
reprecipitated  by  (NH4)2S  ;  but  the  metals  are  thrown  down  as 

insoluble),  they  can  then  be  dissolved  in  water.  The  solution  so  obtained  will 
be  found  to  give  no  precipitate  with  ammonia,  therefore  showing  that  it 
contains  no  barium  phosphate  ;  but  it  will  give  all  the  ordinary  reactions  for 
barium.  Hence  it  follows  that  the  addition  of  phosphoric  acid  to  such  a 
metallic  solution  as  barium  chloride  gives  no  precipitate  of  barium  phosphate. 

*  The  student  will  have  no  difficulty  in  regarding  these  equations  from  the 
ionic  standpoint  without  it  being  necessary  to  dissect  them  further. 

f  The  simpler  formulae  for  the  phosphates  of  Al,  Cr,  and  Fe  will  be  used, 
instead  of  A12(PO4)2,  etc.,  etc. 


66  Qualitative  Analysis. 

sulphides,  while  the  phosphoric  acid  unites  with  the  ammonia, 
and  therefore  remains  in  the  solution  :  thus,  taking  iron  as  an 
example — 


+  2H2S  +  S 
(c)  Precipitation  of  ferric  phosphate  by  BaCO3— 


:  2FeCl8  +  2H3PO4  i    +  3BaCO3  =  3BaCl2  +  2FePO4  +  3CO., 
: ' :  +3H20 

In  the  case  of  solutions  of  calcium  or  magnesium  phosphates, 
the  addition  of  BaCO3  results  in  the  precipitation,  not  of  the  calcium 
or  magnesium  phosphate,  but  of  barium  phosphate  ;  thus — 


(d\     :3MgCL+2H,PO4:   +  3BaCO3  =  3MgCl2  +  Ba3(PO4)2 

3H2O 


II.  By  Acetic  Acid,  H(C2H3O2).—  Of  the  phosphates  of  the 
metals  of  Groups  III.,  IV.,  and  V.,  those  of  aluminium  and  iron 
are  insoluble  in  acetic  acid,  while  chromium  phosphate  is  only 
dissolved  with  difficulty.  The  phosphates  of  the  remaining  metals 
of  these  groups  are  readily  soluble  in  this  acid. 

When  a  phosphate  is  dissolved  in  acetic  acid,  we  may,  for  thef 
same  reasons  as  apply  in  the  case  of  hydrochloric  acid,  regard  the 
solution  as  containing  the  metallic  acetate  and  phosphoric  acid  ; 
thus,  in  the  case  of  calcium  phosphate  — 

Ca3(P04)2  +  6H(C2H302)  =  3Ca(C2H3O2)2  +  2H3PO4* 

Phosphoric  acid^  therefore,  is  incapable  of  giving  a  precipitate 
with  a  solution  of  calcium  acetate  (or  the  acetate  of  any  metal 
whose  phosphate  is  soluble  in  acetic  acid).  It  can,  however,  pre- 
cipitate the  phosphate  from  the  acetate  of  a  metal  whose  phosphate 
is  insoluble  in  acetic  acid.  Thus,  if  phosphoric  acid  is  added  to  a 
solution  of  aluminium  acetate,  aluminium  phosphate  is  thrown 
down,  and  acetic  acid  formed  — 

A1(C2H302)3  +  H3P04  =  A1P04  4-  3H(C2H3O2) 

*  In  this  solution  the  presence  of  the  phosphoric  acid  does  not  prevent  the 
precipitation  of  the  calcium  as  calcium  oxalate  by  the  addition  of  ammonium 
oxalate, 


The  Phosphates.  67 

If  to  a  solution  of  a  phosphate  (e.g.  calcium  phosphate)  in  acetic 
acid,  ferric  chloride  (or  aluminium  chloride)  be  added,  by  double 
decomposition  acetic  acid  is  formed  in  the  solution,  and  ferric  phos- 
phate is  therefore  precipitated  ;  thus — 


;3Ca(C2H302)2+  2H3P04:.   +  2FeCl3  -  2FePO4  +  3CaCl2 

+  6H(C2H302) 

Now,  although  phosphoric  acid  is  unable  to  precipitate  ferric 
phosphate  when  added  to  ferric  chloride  (ferric  phosphate  being 
soluble  in  hydrochloric  acid),  the  presence  of  a  soluble  acetate,  such 
as  sodium  acetate,  would  bring  about  the  same  conditions  as  exist 
in  the  above  reaction,  and  therefore  again  the  phosphate  of  iron  will 
be  thrown  down — 


3Na(C2H302)  +   :  H3PO4  +  FeCl3 :    =  FePO4  +  3NaCl 
: :  +3H(C2H302) 

Perhaps  the  similarity  of  the  conditions  will  be  even  more 
evident  if  the  interacting  compounds  in  the  two  cases  are 
represented  in  the  ionised  state — 


3Ca,:6(C2H302),6H 
3NaJ3(C2H302),3H 


,2P04,2Fej,6Cl 
,P04,      Fe!,3Cl 


i  2*       is    i 

FePO4  precipitated,  while 

acetic    acid    and    sodium    (or   calcium)    chloride   remain   in    the 
solution. 

The  Removal  of  Phosphoric  Acid. — The  phosphoric  acid 
may  be  separated  from  a  phosphate  in  two  ways  :  namely,  either 
by  causing  the  metal  to  form  some  insoluble  compound  with  a 
reagent,  or  by  making  the  phosphoric  acid  unite  with  a  reagent  to 
produce  an  insoluble  compound.  For  example,  the  phosphoric 
acid  may  be  withdrawn  from  the  phosphates  of  iron,  manganese, 
zinc,  cobalt,  and  nickel,  by  means  of  ammonium  sulphide,  as  ex- 
plained on  p.  66.  In  these  cases,  the  metals,  in  the  form  of 
insoluble  sulphides,  are  precipitated,  while  the  phosphoric  acid,  in 
combination  with  ammonia,  remains  in  solution.  The  disadvantage 
of  this  method  lies  in  the  fact  that  the  phosphate  of  ammonia  left 
in  the  liquid,  is  able  to  undergo  double  decomposition  with  any 
other  metallic  salts  which  might  be  present,  and  precipitate  them 


68  Qualitative  Analysis. 

as  phosphates  ;  therefore  the  final  result  would  only  be  that  of  sub- 
stituting one  insoluble  phosphate  for  another.  The  following 
examples  will  make  this  clear  : — 

Suppose  a  mixture  under  examination  to  consist  of  alumina  and 
phosphate  of  iron  ;  on  dissolving  in  hydrochloric  acid,  we  should 
get  the  following  reaction  : — 

A12O3  +  2FePO4  +  I2HC1  =  2A1C13  +  2Fe-Cl3  +  2H3PO4  +  3H2O 

On  adding  ammonium  sulphide  to  the  solution,  the  iron  is  precipi- 
tated as  FeS  and  the  aluminium  as  phosphate  ;  thus— 


2A1C13  +    i2FeCl3  +  2H3P04:   +  6(NH4)2S  =  2FeS  +  2A1PO4 

+  I2NH4C1  +  3H2S  +  S 

Or  again,  if  salts  of  Ba,  Sr,  Ca,  or  Mg  (other  than  phosphates)  are 
present  with  a  phosphate  which  is  decomposed  by  ammonium 
sulphide,  then  the  ammonium  phosphate  which  is  formed  would 
react  upon  such  salts  and  precipitate  the  metals  as  phosphates. 
Thus,  suppose  a  mixture  of  zinc  phosphate  and  calcium  carbonate 
is  being  examined  ;  as  before,  it  is  dissolved  in  hydrochloric  acid, 
and  then,  on  the  addition  of  ammonium  sulphide,  a  precipitate  is 
obtained  consisting  of  a  mixture  of  zinc  sulphide  and  calcium 
phosphate — 

Zn3(P04)2  +  3CaC03  +  I2HC1  -  3ZnCl2  +  2H3PO4 

3H20 


:  3ZnCl2  +  2H3P04 :   +  3CaQ2  +  6(NH4),S  -  3ZnS  +  Ca3(PO4)2 

+  I2NH4C1  +  3H2S  +  S 

In  the  ordinary  course  of  analysis,  on  the  addition  of  NH4C1 
and  NH4HO  for  the  precipitation  of  Group  IIlA.,  if  phosphates  are 
present  they  are  also  precipitated  along  with  the  hydroxides  of  Al, 
Cr,  and  Fe.  If,  therefore,  ammonium  sulphide  be  added,  any 
phosphates  of  Fe,  Zn,  Mn,  Ni,  and  Co  which  are  present  will  be 
converted  into  sulphides,  and  the  ammonium  phosphate  which  is 
formed  will  react  upon  any  compounds  of  Ba,  Sr,  Ca,  and  Mg  which 
might  be  present,  not  as  phosphates,  and  convert  them  into  phos- 
phates. In  analysis,  therefore,  the  second  method  for  removing 
phosphoric  acid  is  had  recourse  to,  namely,  that  of  employing  a 
reagent  which  will  form  an  insoluble  compound  with  the  acid,  and 
so  throw  it  out  of  solution. 


Separation  of  Phosphoric  Acid.  69 

The  reagent  employed  is  ferric  chloride,  and  the  separation  is 
based  upon  the  fact  already  explained,  viz.  that  when  this  salt  is 
added  to  an  acetic  acid  solution  of  a  phosphate,  it  precipitates  the 
phosphoric  acid  as  ferric  phosphate,  leaving  the  metal,  which  was 
originally  united  to  the  phosphoric  acid,  in  the  solution  as  a 
chloride.  The  separation  is  made  in  the  following  manner  : — 


SEPARATION  OF  PHOSPHORIC  ACID. 

The  precipitate  produced  by  NH4HO  in  presence  of  NH4C1 
(which  may  contain  the  hydroxides  of  Al,  Cr,  and  Fe,  as  well  as  all 
the  phosphates),  is  dissolved  in  a  little  warm  dilute  HC1,  and  the 
solution  made  nearly  neutral  by  the  addition  of  Na2CO3.*  A  mixture 
of  sodium  acetate  and  acetic  acid  t  is  then  added,  and  the  solution 
boiled  and  filtered.? 

*  This  is  in  order  to  remove  the  excess  of  HC1  used  in  dissolving  the  pre- 
cipitate. This  acid  would,  in  any  case,  be  neutralised  by  double  decomposi- 
tion with  the  sodium  acetate,  when  that  reagent  is  added  ;  but  it  is  better  to 
get  rid  of  it  previously  by  means  of  sodium  carbonate,  so  that  the  sodium 
acetate  may  be  utilised  entirely  in  bringing  about  the  reactions  explained  on 
p.  67,  and  also  in  the  footnote  below. 

t  See  list  of  reagents  (Appendix). 

%  If  the  whole  of  the  phosphoric  acid  contained  in  the  substance  under- 
going analysis  was  originally  present  in  combination  with  Al,  Cr,  or  Fe,  then 
by  this  reaction  the  whole  of  the  phosphoric  acid  will  have  been  thrown  out  of 
solution.  Moreover,  even  if  some  of  the  phosphoric  acid  had  been  originally 
present  in  combination  with  other  metals,  it  might  happen  that  there  was  also 
present  enough  aluminium  or  iron,  not  as  phosphate,  to  combine  with  all  the 
phosphoric  acid  of  these  other  phosphates,  and  so  to  cause  the  complete 
removal  of  this  acid.  The  following  example  will  make  this  clear  :  Suppose 
a  mixture  of  aluminium  phosphate,  ferric  oxide,  and  calcium  phosphate  is  to  be 
examined.  When  this  has  been  dissolved  in  HC1,  we  may  represent  the  action 
of  sodium  acetate  in  the  following  manner  : — 


;  A1C13  +  H3P04  •  +  2FeCl3  +    j  3CaCl2  +  2H3PO4  |  +  9Na(C2H3O2) 

^ — 
\  }   precipitated 

'    sHfwM      +          6H(Qa£o2    +    3CaCl,      }  in  solution 

If  more  than  enough  ferric  chloride  is  present  than  is  required  to  take  the 
whole  of  the  phosphoric  acid  from  the  calcium  phosphate  (or  other  similar 
phosphates),  then  all  the  phosphoric  acid  is  thrown  down,  and  the  excess  of 
the  FeCl3  passes  into  the  solution.  But  if,  on  the  other  hand,  there  is  an 
excess  of  calcium  phosphate,  it  passes  into  solution  owing  to  the  free  acetic 
acid  present.  If  no  compounds  of  Al,  Cr,  or  Fe  are  present,  either  as  phos- 
phates or  otherwise,  no  precipitate  will  be  produced  by  the  sodium  acetate. 


Qualitative  A  nalysis. 


The  precipitate  con- 
tains phosphates  of 
Al,  Cr,  or  Fe. 

This  precipitate  may  be 
treated  with  Na2O2 
exactly  as  the  hy- 
droxides, p.  49.  A1PO4 
dissolves  in  the  NaHO, 
and  CrPO4  js  oxidised 
intoNa2CrO4.  On  filter- 
ing, the  iron  is  left  on 
the  filter.  The  Al  may 
be  detected  by  neutra- 
lisation and  reprecipita- 
tionwithNH4HO.  The 
chromium  by  BaCl2 
in  acetic  acid  solution. 
(Barium  phosphate  is 
not  precipitated  in 
presence  of  acetic  acid. ) 


The  filtrate.  FeCl3  is  added  drop  by  drop  *  until 
precipitation  is  complete,  at  which  point  the 
liquid  will  begin  to  assume  a  distinct  reddish 
colour  (due  to  the  formation  of  soluble  ferric 
acetate). t  The  mixture  is  gently  boiled  for  a 
few  minutes  (whereby  the  ferric  acetate  is  con- 
verted into  an  insoluble  basic  acetate),  and  then 
filtered. 

The  precipitate,  containing  the  whole  of  the 
phosphoric  acid  as  ferric  phosphate  and 
more  or  less  basic  ferric  acetate,  is  thrown 
away.  The  solution  now  contains  the  metals 
of  Groups  Ills,  and  IV.,  as  well  as  Mg,  in 
the  form  of  chlorides.  Their  separation  is 
made  by  the  usual  methods. 


APPENDIX    TO    CHAPTER    VII. 

THE    RARE   METALS   OF   GROUP    Ill-t 

These  elements  may  for  convenience  be  arranged  in  five  groups, 
based  upon  the  composition  of  the  precipitates  which  are  thrown 
down  by  the  group-reagents  : 

*  If  the  first  drops  of  FeCl3  produce  no  precipitate,  it  proves  that  all  the 
phosphoric  acid  has  already  been  thrown  down  (see  previous  note).  The 
metals  are  all  present,  therefore,  as  chlorides,  and  their  separation  may  be 
made  by  the  ordinary  methods. 

f  Ferric  acetate  is  formed  by  the  action  of  the  excess  of  ferric  chloride  upon 
the  sodium  acetate  present ;  thus — 

FeCl3  +  3Na(C2H302)  =  Fe(C2H3O2)3  +  3NaCl 

It  is  very  important  to  avoid  the  use  of  any  unnecessary  excess  of  FeCl3>  for  the 
precipitated  ferric  phosphate  is  soluble  both  in  ferric  chloride  and  in  ferric 
acetate.  The  transformation  of  soluble  ferric  acetate  into  the  insoluble  basic 
acetate  by  boiling,  is  expressed  as  follows  : — 

3Fe(C2H302)3  +  3H20  =  Fe(C2H3O2)3,Fe2O3  +  6H(C2H3O2) 

In  cases  where  very  small  quantities  of  a  phosphate  and  a  great  excess  of  iron 
are  present  in  the  substance  under  examination,  it  is  necessary  to  reduce  the 
iron  to  the  "  ferrous  "  state  by  means  of  sulphurous  acid.  Ferrous  salts  do  not 
dissolve  ferric  phosphate. 

J  A  systematic  study  of  the  reactions  of  the  rare  metals  lies  entirely  outside 
the  scope  of  this  book.  In  many  cases  it  is  almost  impossible  to  purchase  the 
compounds  in  anything  approaching  to  a  condition  of  purity,  while  with  most 
of  the  metals  the  cost  of  the  salts  practically  prohibits  such  a  course  (see  note, 
p.  26).  These  elements,  however,  are  not  all  equally  "  rare,"  and  therefore  in 


The  Rare  Metals  of  Group  III. 


(I) 

(2) 

(3) 

(4) 

(5) 

Beryllium 

Scandium 

Zirconium 

Titanium 

Uranium 

Yttrium 

Thorium 

Tantalum 

Indium 

Ytterbium 

Niobium 

(Thallium)* 

Cerium 

Vanadium  t 

Lanthanum 

Those  of  the  first  three  divisions  are  precipitated  in  the  form  of 
hydroxides,  which  are  basic  in  their  character.  Those  of  the  fourth 
division  are  thrown  down  in  the  form  of  hydrated  oxides  of  a  more 
or  less  acidic  character,  while  the  remaining  metals  come  down  as 
sulphides.  The  composition  of  the  various  compounds  is  as 
follows  : — 


(I) 

Be(HO), 

(2) 

Sc(HO)3 
Y(HO)3 
Yb(HO)3 
Ce(HO)3 
La(HO)3 

(3) 
Zr(HO)4 
Th(HO)4 

(4) 
H2TiO3 
H3TaO4 
H3NbO4 

(5) 
(U02)S 
InS 
(T12S)2 

Beryllium. — Named  from  beryl,  the  chief  source  of  the  ele- 
ment. The  sulphate,  BeSO4,  is  the  most  readily  obtained  salt. 

The  reactions  of  this  element  resemble  very  closely  those  of 
aluminium — so  much  so,  that  at  one  time  the  two  elements  were 
regarded  as  belonging  to  the  same  natural  family.  Thus,  ammonia, 
potash,  and  soda  produce,  with  solutions  of  either  metal,  a  white 
flocculent  precipitate  of  the  respective  hydroxides,  A12(HO)6  and 
Be(HOV 

Both  precipitates  are  soluble  in  excess  of  the  fixed  alkalies,  but 
the  beryllium  compound  is  reprecipitated  if  the  diluted  solution  is 
boiled. 

Beryllium  hydroxide  is  decomposed  by  boiling  with  ammonium 
chloride,  ammonia  being  evolved  and  beryllium  chloride  passing 
into  solution  ;  thus —  ' 

Be(HO)2  +  2NH4C1  =  BeCl2  +  2NH3  +  2H2O 

Aluminium  hydroxide  undergoes  no  change  when  similarly  treated, 
hence  the  metals  may  be  separated  by  this  reaction. 

The  carbonates  of  the  alkalies  give  precipitates  of  beryllium 
carbonate,  or  basic  carbonate  ;  readily  soluble  in  excess  of  ammo- 
nium carbonate,  giving  a  double  carbonate. 

The  precipitate  is  soluble,  but  far  less  readily,  in  the  carbonates 
of  the  fixed  alkalies.  Beryllium  is  therefore  readily  separated  from 

this  section  some  of  the  characteristic  reactions  of  a  few  of  the  most  commonly 
occurring  of  these  metals  are  given. 

*  Thallium  also  appears  in  Group  I.  (see  p.  16),  since  the  chloride  is 
precipitated  by  hydrochloric  acid  from  thallous  solutions  ;  and,  like  lead 
chloride,  is  slightly  soluble  in  water. 

f  See  footnote  on  p.  18. 


72  Qualitative  Analysis. 

aluminium  by  adding  ammonium  carbonate  and  warming  the 
solution.  Aluminium  hydroxide  is  precipitated,  while  the  beryllium 
goes  into  solution  as  the  double  carbonate  of  ammonium  and 
beryllium.  On  filtering  off  the  aluminium  hydroxide,  the  beryllium 
hydroxide  may  be  precipitated  by  ammonia,  after  first  neutralising 
with  hydrochloric  acid. 

The  salts  of  beryllium  possess  a  characteristic  sweet  taste, 
hence  the  name  glucinum,  which  was  formerly  applied  to  the 
element. 

Zirconium. — The  oxide  of  this  element,  along  with  others  of 
the  so-called  "  rare  earths  "  (but  more  especially  Zirconia),  has  the 
property  of  remaining  unchanged  for  a  long  time  when  heated  to 
incandescence,  and  of  emitting  a  bright  white  light  when  so  heated. 
It  is  on  this  account  used,  with  others  of  the  rare  earths,  in  the 
construction  of  the  "  mantles  "  of  the  incandescent  gas-burners 
now  so  common. 

Alkaline  hydroxides,  as  well  as  ammonium  sulphide  (group- 
reagent),  give  a  white  precipitate  of  zirconium  hydroxide,  Zr(HO)4, 
resembling  aluminium  hydroxide  in  appearance.  It  is  distinguished 
from  the  latter  in  that  it  is  insoluble  in  excess  of  potassium  or 
sodium  hydroxide. 

Alkaline  carbonates  give  a  white  precipitate  consisting  of  a 
basic  carbonate,  which  is  soluble  in  excess,  especially  of  ammonium 
carbonate. 

Potassium  sulphate  (a  concentrated  boiling  solution)  gives  a 
precipitate  of  a  double  sulphate,  which,  when  thrown  down  from 
the  hot  solution,  is  scarcely  soluble  in  hydrochloric  acid  (thorium 
gives  a  similar  precipitate  under  the  same  conditions,  but  the 
thorium  compound  is  soluble  in  hydrochloric  acid).  In  dilute 
solutions  the  precipitate  only  appears  after  standing  for  some  hours. 
Sodium  thiosulphate  gives  a  precipitate  of  zirconium  thiosulphate, 
which  on  boiling  is  complete  even  in  very  dilute  solutions  (thorium 
behaves  similarly).  Oxalic  acid  gives  a  white  precipitate  of 
zirconium  oxalate,  soluble  in  ammonium  oxalate.  (Thorium  oxalate, 
similarly  precipitated,  is  insoluble  in  ammonium  oxalate.) 

Hydrogen  peroxide,  added  to  a  slightly  acid  solution  of  a 
zirconium  salt,  gives  a  white  precipitate,  believed  to  be  either  ZrO3 
or  Zr2O5.  (Niobium  and  titanium  do  not  give  a  precipitate.) 

By  means  of  these  three  reactions  zirconium  can  be  separated 
from  the  other  "rare  earths  "  of  this  group. 

The  hydrated  oxides,  which  are  precipitated  by  ammonia,  are 
first  dissolved  in  hydrochloric  acid,  and  the  solution  nearly 
neutralised  with  sodium  carbonate.  Sodium  thiosulphate  is  then 
added,  and  the  mixture  boiled.  The  precipitate  may  consist  of  the 
thiosulphates  of  zirconium  and  thorium,  together  with  titanic  acid. 
The  precipitate  is  treated  with  boiling  hydrochloric  acid,  which 
dissolves  the  two  thiosulphates,  and  possibly  a  little  titanic  acid. 
Excess  of  ammonium  oxalate  is  added  to  the  solution,  which  at 
first  precipitates  oxalates  of  zirconium  and  thorium,  but  redissolves 
the  zirconium  oxalate.  The  insoluble  thorium  oxalate  is  removed 


The  Rare  Metals  of  Groiip  III.  73 

by  filtration,  and  any  titanium  in  the  solution  is  precipitated  by 
means  of  ammonium  carbonate,  added  in  excess  in  order  to  re- 
dissolve  the  zirconium  basic  salt,  which  is  first  thrown  down. 
Any  permanent  precipitate  is  filtered  off,  and  the  filtrate  con- 
centrated by  gentle  evaporation.  A  boiling  strong  solution  of 
potassium  sulphate  is  then  added,  which  precipitates  the  double 
sulphate  of  potassium  and  zirconium.* 

Titanium. — This  element  is  met  with  in  small  quantities  in 
many  specimens  of  iron  ores,  clays,  and  igneous  rocks.  The 
minerals  special  to  the  element,  rutile,  anatase  (TiO2) ;  sphene 
(silicate  and  titanate  of  calcium),  are  rare  substances — titantferous 
iron  (ferrous  titanate)  is  less  rare. 

Titanium  oxide,  TiO2,  being  insoluble  in  hydrochloric  and 
nitric  acid,  is  found  in  the  insoluble  residue  after  treatment  with 
acids ;  titanates,  on  the  other  hand,  are  dissolved  by  hydrochloric 
acid,  but  on  boiling  the  solution  white  titanic  acid,  H2TiO3,  is 
precipitated. 

Titanium  oxide  (in  the  absence  of  other  metals  which  colour  a 
bead  of  microcosmic  salt),  when  heated  in  a  bead  of  microcosmic 
salt  in  the  inner  blowpipe  flame,  imparts  to  the  bead  a  colour 
which  is  yellowish  when  hot,  but  which  becomes  violet  as  the  bead 
cools.  The  colour  is  more  readily  obtained  by  the  aid  of  some 
additional  reducing  agent  besides  the  reducing  flame  ;  thus  if  the 
salt  be  heated  on  charcoal  instead  of  a  platinum  wire,  or  if  a  trace 
of  zinc  be  added,  the  result  is  more  quickly  obtained.  In  the 
presence  of  small  quantities  of  iron  the  bead  appears  brown-red. 

Titanium  dioxide  is  separated  from  silicon  dioxide  (or  silicates) 
by  the  action  of  hydrofluoric  acid  (sulphuric  acid  being  present  to 
prevent  the  volatilisation  of  titanium,  as  fluoride).  Titanium  is 
obtained  in  solution  (and  separated  from  silica,  and  also  from 
compounds  of  tantalum  and  niobium)  by  fusion  with  hydrogen 
potassium  sulphate.  The  "  melt  "  is  extracted  with  cold  water, 
when  the  titanium  passes  into  solution  as  a  sulphate. 

When  titanium  oxide  is  fused  with  potassium  carbonate, 
potassium  titanate,  K2TiO3,  is  formed,  which,  being  insoluble  in 
water,  may  be  separated  from  any  alkaline  silicate  by  extracting 
the  melt  in  cold  water. 

If  the  residue  of  potassium  titanate  be  then  treated  with  cold 
dilute  hydrochloric  acid,  it  dissolves,  the  solution  containing  titanic 
acid,  Ti(HO)4  or  TiO2,2H2O.  If  this  solution  be  heated,  the 

*  This  separation,  and  the  reactions  for  zirconium,  may  be  made  by  dis- 
solving up  a  couple  of  the  incandescent  gas  "  mantles  ;  "  either  new  ones,  or 
those  which  have  become  worn  out  by  use.  In  the  former  case  they  should 
first  be  set  fire  to,  in  order  to  burn  off  the  organic  matter  present.  The 
mantles  are  boiled  in  a  test-tube  with  a  little  strong  sulphuric  acid,  and  the 
liquid  when  cold  diluted  with  water.  The  insoluble  residue  (consisting  of 
the  main  portion  of  the  material)  is  filtered  off,  and  the  hydrated  oxides  of  the 
rare  earths,  along  with  any  alumina,  are  precipitated  by  the  addition  of 
ammonia.  The  precipitate  is  then  washed,  and  dissolved  in  a  small  quantity 
of  hydrochloric  acid. 


74  Qualitative  Analysis. 

titanic  acid  is  rendered  insoluble,  and  is  therefore  precipitated, 
the  precipitate  being  the  hydrated  oxide  H2TiO3,  or  TiO(HO)2, 
or  TiO2,H2O. 

If  to  the  cold  solution  of  titanic  acid  there  be  added  ammonia, 
or  the  hydroxide  or  carbonate  of  the  alkalies,  or  ammonium 
sulphide,  a  white  precipitate  is  produced,  consisting  of  the  hydrated 
compound  Ti(HO)4,  or  TiO2,2H2O,  which  redissolves  readily  in 
dilute  hydrochloric  acid  or  sulphuric  acid.  Moderate  rise  of 
temperature  at  once  converts  this  soluble  compound  into  the 
insoluble  titanic  acid,  H2TiO3,  or  TiO2,H2O. 

Uranium. — Like  the  element  chromium,  uranium  forms 
uranous  and  uramV  salts,  as  well  as  manates. 

The  uranous  salts  are  derived  from  uranous  oxide,  UO2,  in 
which  the  element  is  tetravalent.  Thus,  uranous  sulphate  is 
represented  by  the  formula  U(SO4)2. 

These  salts  readily  pass  by  oxidation  into  uranic  compounds, 
and  are  therefore  powerful  reducing  agents.  They  are  for  the 
most  part  green  in  colour.  The  uranic  or  uranyl  salts  are  derived 
from  the  oxide,  UO3,  or  (UO2)O,  in  which  the  element  is  hexa- 
valent.  They  are  regarded  as  containing  the  divalent  radical 
uranyl,  (UO2),  which  takes  the  place  of  a  divalent  metal ;  thus, 
uranyl  sulphate  and  nitrate  are  expressed  by  the  formulae 
(UO2)SO4,3H2O  and  (UO2)(NO3)2,6H2O  respectively.  These 
salts  are  mostly  yellow  in  colour,  and  soluble  in  water. 

The  uranates  are  constituted  like  the  dichromates  (uranates 
corresponding  to  normal  chromates  are  not  known).  Thus,  sodium 
uranate,  Na2U2O7,  analogous  to  sodium  dichromate,  Na2Cr2O7. 

Uranyl  nitrate  and  acetate  are  the  salts  most  commonly  met 
with,  being  used  in  the  volumetric  determination  of  phosphoric 
acid. 

Alkaline  carbonates  give  with  uranyl  salts  a  yellow  precipitate, 
consisting  of  a  double  carbonate  of  uranium  and  the  alkali.  The 
precipitate  is  soluble  in  excess  of  the  reagent. 

The  uranium  in  this  solution  is  not  precipitated  by  ammonium 
sulphide.  If  the  solution  be  neutralised  with  acid,  a  uranate  of 
the  alkali  is  thrown  down. 

With  uranous  salts,  these  reagents  give  a  green  precipitate, 
also  soluble  in  excess. 

Caustic  alkalies  give  with  tiranyl  salts  a  yellow  precipitate  of 
the  uranate  of  the  alkali,  insoluble  in  excess  of  the  reagent. 

With  uranous  salts,  these  reagents  give  a  chocolate-coloured 
precipitate  of  uranous  hydroxide,  U(HO)4. 

Ammonium  sulphide  gives  with  uranyl  salts  a  brown  precipitate 
of  uranyl  sulphide,  (UO2)S,  insoluble  in  excess  of  the  reagent,  but 
soluble  in  normal  ammonium  carbonate. 

In  acid  solutions  sulphuretted  hydrogen  reduces  uranyl  to 
uranous  compounds. 

Potassium  ferrocyanide  gives  with  uranyl  salts  a  brown  pre- 
cipitate. This  reagent  is  employed  as  an  indicator  in  the  volu- 
metric estimation  of  phosphoric  acid  by  means  of  uranyl  salts. 


CHAPTER   VIII. 
THE  METALS  OF  GROUP  II. 

THE  metals  of  this  group  (as  well  as  those  of  Group  I.)  are 
characterised  by  their  common  property  of  forming  sulphides  in  an 
acid  solution  ;  that  is  to  say,  their  sulphides  are  insoluble  in  dilute 
acids.  By  this  property  they  are  all  sharply  separated  from  the 
metals  of  Group  III.  The  acid  which  has  been  found  to  be  the 
most  convenient  to  have  presentr  is  hydrochloric  acid  ;  and  since, 
on  acidifying  with  this  acid  preparatory  to  the  precipitation  of  the 
sulphides,  the  chlorides  of  silver,  lead,  and  mercurous  mercury  *  are 

*  Mercury  forms  two  classes  of  salts,  mercuric  and  mercurous ;  and  the 
relation  in  which  they  stand  to  each  other  is  very  interesting.  In  the  com- 
pounds of  the  first  type,  the  metal  is  playing  the  part  of  an  ordinary  divalent 
element,  replacing  two  atoms  of  hydrogen  in  acids  and  giving  such  salts  as 
HgCl2,  Hg(NO3)2,  KgSO4,  HgS,  etc.  In  the  mercurous  compounds  the  pro- 
portion of  mercury  to  the  negative  radical  is  twice  as  great  as  in  the  mercuric 
salts,  hence  the  composition  of,  say,  mercurous  chloride  may  be  expressed 
either  by  the  formula  Hg2Cl2  or  HgCl.  Some  chemists  adopt  the  latter 
formula,  and,  regarding  the  mercury  in  mercurous  compounds  as  acting  the 
part  of  a  monovalent  element,  express  the  various  salts  by  such  formulae  as 
HgNO3,  Hg2SO4,  etc.  Others  prefer  to  consider  the  mercurous  salts  as  com- 
pounds, in  which  the  divalent  radical,  or  double  atom,  (Hg2),  is  substituted  for 
the  single  divalent  atom,  (Hg),  and  therefore  express  the  compounds  by  such 
formulas  as  Hg2Cl2,  Hga(.NOj)j,  Hg2SO4.  The  density  of  the  vapour  yielded 
by  heating  mercurous  chloride  is  117  '59,  which,  being  half  that  demanded  by 
the  formula  Hg2Cl2,  gave  support  to  the  view  that  HgCl  was  the  correct 
formula.  But  it  has  since  been  shown  that  tke  compound  dissociates  on 
heating,  into  mercuric  chloride,  HgCl2,  and  Hg  (mercury  giving  monatomic 

molecules).  On  solution  in  water,  the  mercuric  salts  yield  the  divalent  ion  Hg, 
with  its  two-unit  electric  charges,  while  the  mercurous  compounds  may  be 
regarded  as  furnishing  two  monovalent  ions  Hg,  Hg,  each  with  its  one  charge. 
In  this  book  the  formulae  adoped  for  mercurous  compounds  will  be  those 
which  represent  them  as  containing  the  double  atom  (Hg2).  As  the  two 
classes  of  compounds  present  great  differences  in  their  chemical  reactions, 
behaving,  indeed,  more  like  compounds  of  two  different  metals,  one  of  which 
belongs  to  Group  I.  and  the  other  to  Group  II.,  there  is  at  least  some 
advantage  in  employing  a  symbol  for  mercurous  mercury  (Hg2),  which  at  a. 
glance  distinguishes  it  from  that  used  to  denote  mercuric  mercury,  Hg.  All 
the  salts  of  mercury  will  therefore  be  formulated  as  salts  of  a  dibasic  metallic 
radical,  which  in  the  mercuric  compounds  has  the  symbol  Hg,  and  in  the 
mercurous  salts  (Hg2)  (with  or  without  the  bracket).  Thus — 

Chlorides,  HgCl,        (Hg2)Cl2 

Sulphates,  HgSO4       (Hg2)SO4 


76  Qualitative  Analysis. 

thrown  down,  these  three  metals  are  treated  separately,  and  con- 
stitute Group  I. 

The  precipitation  of  these  three  metals  as  chlorides,  however, 
is  only  complete  in  the  case  of  Ag  and  (Hg2).  Lead  chloride  is 
soluble  to  some  extent  even  in  cold  water,  hence  a  portion  of  the 
lead  passes  through  into  Group  II. 

The  metals  of  Group  II.  are  divided  into  two  sections,  namely  — 

(1)  Metals     whose    sulphides     are    insoluble    in     ammonium 
sulphide  — 

Mercury  (mercuric),  lead,  bismuth,  cadmium,  copper. 

(2)  Metals  whose  sulphides  dissolve  in  ammonium  sulphide- 

Arsenic,  antimony,  tin,  gold,  platinum. 

REACTIONS  OF  THE  METALS  OF  GROUP  II.—  DIVISION  i. 
Mercury,  Hg. 

DRY  REACTIONS.—  When  heated  alone  in  a  tube,  many  mercury 
compounds  (those  with  the  halogens,  for  example)  volatilise  un- 
changed, giving  sublimates  of  the  same  compound.  The  iodide 
(red)  when  heated  forms  a  sublimate,  consisting  chiefly  of  the  yellow 
allotropic  form  of  HgI2,  which  when  cold  changes  to  red  if  scratched 
or  rubbed.  Some  mercury  compounds  when  heated  decompose, 
and  metallic  mercury  volatilises  and  sublimes  in  the  tube. 

If  a  mercury  salt  be  mixed  with  several  times  its  weight  of 
sodium  carbonate  (both  being  as  dry  as  possible),  and  the  mixture 
be  strongly  heated  in  a  dry  narrow  test-tube,  a  sublimate  of  metallic 
mercury  will  be  obtained.  The  sublimed  mercury  will  present  the 
appearance  of  a  bright  metallic  mirror,  but  if  examined  by  means 
of  a  lens,  or  if  rubbed  with  a  glass  rod,  distinct  globules  of  liquid 
metal  will  be  visible. 

WET  REACTIONS.—  (a)  Mercuric  Compounds.—  Of  the  com- 
mon salts,  the  nitrate,  sulphate,  chloride,  and  bromide  (but  not  the 
iodide)  are  soluble  in  water,  but  the  solubility  is  not  very  great. 

KHO  or  NaHO  gives  with  mercuric  compounds  a  yellow 
precipitate  *  of  mercuric  oxide,  HgO  — 

HgCl2  +  2KHO  =  HgO  +  H20  +  2KC1 


_ 

Double  ammonium  compounds,  NH2HgCl     N 

*  On  the  first  addition  of  the  reagent,  the  precipitate  appears  a  browmsh 

t^^^ 
«aSKtajion  should  be  yellow,  J^  ttot  PT  epared  »  the 

dry  way  is  brick-red,  is  not  known.     Compare  also  the  sulphide. 


Group  II. — Division  i.  77 

The  precipitate  is  insoluble  in  excess  of  the  reagent. 

NH 4HO  produces  a  white  precipitate  of  an  ammoniacal  mercuric 
compound,  where  two  atoms  of  hydrogen  from  the  ammonium 
radical  are  replaced  by  the  divalent  atom  Hg  ;  thus — 

HgCl2  +  2NH4HO  =  NH2HgCl  +  NH4C1  +  2H2O 

Or  with  mercuric  nitrate — 

Hg(N03)2  +  2NH4HO  =  (NH2Hg)N03  +  NH4NO3  +  2H2O 

H2S  produces  a  black  *  precipitate  of  HgS.  The  precipitation 
is  only  complete  after  some  time,  and  when  the  solution  is  con- 
siderably dilute.  The  compound  is  insoluble  in  HC1,  and  in 
HNO3  even  when  boiling.  (The  prolonged  action  of  boiling  HNO3 
partially  converts  it  into  the  white  compound  Hg(NO3)2,  2HgS.) 
Mercuric  sulphide  dissolves  in  aqua  regia,  forming  mercuric 
chloride.  In  the  presence  of  caustic  alkalies  it  dissolves  in  sodium 
or  potassium  sulphide  (not  in  ammonium  sulphide\  forming  the 
double  sulphides,  HgS,Na2S  and  HgS,K2S. 

(NH4)2S  gives  the  same  precipitate.  The  same  result  is  also 
obtained  by  the  addition  of  sodium  thiosulphate,  Na2S2O3,  to  a  warm 
solution  acidified  with  HC1. 

El  precipitates  HgI2  as  a  rich  scarlet  compound,  soluble  in 
excess  of  either  solution.  When  first  precipitated  it  appears 
yellow,  but  quickly  turns  salmon-red  and  then  scarlet.  The  com- 
pound is  dimorphous,  and  can  be  obtained  either  in  the  red 
(quadratic  crystals)  or  the  yellow  (rhombic  prisms]  variety. 

Reduction  of  Mercuric  Compounds. — By  reducing  agents 
mercuric  compounds  may  be  converted  into  mercurous  salts,  or  the 
reduction  may  go  a  stage  further  and  result  in  the  precipitation  of 
mercury  in  the  metallic  state.  Thus,  on  the  addition  of  stannous 

*  The  progress  of  this  precipitation  is  accompanied  by  very  characteristic 
changes  of  colour.  The  first  action  of  the  H2S  is  to  give  a  white  precipitate, 
which  then  passes  through  various  shades  of  colour,  from  yellow  to  yellowish- 
red,  to  brown,  and  lastly,  black.  The  white  substance  is  a  compound  of  HgS 
with  the  mercuric  salt  in  solution,  HgCl2,2HgS,  or  Hg(NO3)2  .aHgS,  and  the 
changes  in  colour  are  ascribed  to  the  gradual  conversion  of  this  into  black,  HgS. 
It  does  not  appear  quite  obvious,  however,  how  a  gradual  alteration  in  the  pro- 
portions of  the  -white  double  compound  and  the  black  sulphide  can  give  the 
orange  and  reddish  tints  which  are  seen.  Mercuric  sulphide  prepared  by  other 
processes  is  red  (the  pigment  known  as  vermilion}.  Why  the  compounds  pre- 
pared in  different  ways  should  be  so  very  different  in  colour  is  not  known  ; 
probably  it  is  a  case  of  dimorphism  similar  to  that  exhibited  by  HgI2,  and  it 
may  be  that  to  some  extent  the  red  HgS  is  precipitated,  but  is  not  stable 
under  the  conditions  which  are  present,  and  so  passes  into  the  black 
modification. 


78  Qualitative  Analysis. 

chloride,  SnCl2,  a  white  precipitate  of  mercurous  chloride  is 
produced — 

2HgCl2  +  SnClo  =  (Hg2)Cl2  +  SnCl4 

On  gently  warming  with  an  excess  of  stannous  chloride,  the  pre- 
cipitated mercurous  chloride  changes  to  a  grey  deposit  of  mercury 
in  a  condition  of  fine  powder — 

(Hg)2Cl2  +  SnCl2  =  SnCl4  +  2Hg 

Many  metals  are  capable  of  displacing  mercury  from  its  solutions, 
the  mercury  being  deposited  upon  the  metal.  Thus,  if  a  strip  of 
clean  copper  be  immersed  in  a  neutral  or  slightly  acid  solution  of  a 
mercury  salt,  it  becomes  coated  with  a  white  silvery  deposit  of  an 
amalgam  of  copper  and  mercury,  from  which  the  mercury  can  be 
readily  volatilised  and  obtained  as  a  metallic  sublimate  by  heating 
the  copper  in  a  dry  test-tube.  In  the  case  of  mercuric  salts,  the 
action  may  be  regarded  as  taking  place  in  two  stages  ;  thus — 

2Hg(N03)2  +  Cu  =  (Hg2)(N03)2  +  Cu(N03)2 
(Hg2)(N03)2  +  Cu  =  2Hg  +  Cu(N03)2 

(<£)  Mercurous  Compounds. — Of  the  common  salts  mer- 
curous nitrate  is  the  only  one  which  is  readily  soluble,  and  this  only 
so  long  as  the  water  is  acid  with  nitric  acid.  The  addition  of  much 
water  results  in  the  precipitation  of  a  basic  nitrate.  Mercurous 
sulphate  is  soluble  with  difficulty. 

KHO  or  NaHO  throws  down  a  black  precipitate  of  (Hg2)O. 
Mercurous  oxide  is  very  unstable.  When  gently  warmed,  or  even 
upon  exposure  to  light,  it  is  converted  into  HgO  and  Hg. 

NH  HO  precipitates  an  ammoniacal  mercurous  compound, 
which  is  black.  Its  composition  is  exactly  analogous  to  the 
corresponding  mercuric  compound — 

Hg2(N03)2  +  2NH4HO  =  NH2(Hg2)N03  +  NH4NO3 

H2S  produces  a  black  precipitate,  which  is  a  mixture  of  HgS 
and  Hg.  (Hg2S  is  not  known  to  exist.)  This  precipitate,  therefore, 
behaves,  on  treatment  with  nitric  acid,  in  the  same  way  as  that 
obtained  from  a  mercuric  solution,  If  we  imagine  the  free  atom 
of  mercury  as  first  dissolving  in  the  acid,  the  mercuric  nitrate  so 
formed  unites  with  HgS,  giving  the  white  insoluble  compound 
Hg(N03)2,2HgS. 

(NH4)2S  gives  the  same  precipitate,  but  in  this  case  the  free 
mercury  will  be  also  converted  into  HgS  in  proportion  as  the 
ammonium  sulphide  contains  more  or  less  polysulphide,  for  alkaline 
polysulphides  convert  metallic  mercury  into  HgS, 


Group  II, — Division  I.  79 

HC1  and  soluble  chlorides  precipitate  white  mercurous  chloride, 
Hg2Cl2.  Insoluble  in  dilute  acids  ;  soluble  in  boiling  HNO3,  being 
converted  into  HgCl2  and  Hg,  and  the  mercury  then  dissolves  to 
mercuric  nitrate,  with  evolution  of  oxides  of  nitrogen. 

Long  boiling  with  concentrated  HC1  decomposes  Hg2Cl2  into 
HgCl2  (which  dissolves)  and  Hg,  which  separates  (mercury  being 
insoluble  in  HC1).  Chlorine  water  converts  it  into  mercuric 
chloride  ;  thus — 

Hg2Cl2  +  C12  =  2HgCl2 

Ammonia  converts  it  into  black  mercurous  ammonium  chloride, 
NH2(Hg2)Cl.  (This  constitutes  one  of  the  most  characteristic 
reactions  for  mercurous  compounds.) 

Mercurous  salts  are  reduced  to  metallic  mercury  by  the  reducing 
agents  which  reduce  the  mercuric  compounds  ;  thus,  with  stannous 
chloride  a  grey  precipitate  of  mercury  is  at  once  produced — 

Hg2(NO3)2  +  SnCl2  +  2HC1  =  SnCl4  =  2HNO3  +  2  Hg 

Lead,  Fb. 

DRY  REACTIONS. — Lead  compounds  are  very  readily  reduced 
when  heated  upon  charcoal  before  the  blowpipe  flame,  either  alone 
or  mixed  with  sodium  carbonate  or  potassium  cyanide.  Globules 
of  metallic  lead  are  thus  obtained,  and  at  the  same  time  a  yellowish 
incrustation  is  formed,  consisting  of  the  oxide,  PbO  (litharge]. 
When  cold,  one  of  the  globules  can  be  removed  and  the  properties 
of  the  metal  examined.  Lead  may  be  recognised  by  its  malleability 
and  softness,  the  latter  property  enabling  it  to  leave  a  black  mark 
when  rubbed  upon  paper.  It  is  insoluble  in  cold  HC1  or  H2SO4, 
but  readily  dissolves  in  HNO3,  forming  Pb(NO3)2,  which,  being 
insoluble  in  nitric  acid,  remains  as  a  white  deposit,  but  which 
dissolves  on  dilution  with  water. 

WET  REACTIONS. — The  only  salts  of  lead  which  are  met  with 
in  analysis  are  derived  from  plumbic  oxide,  PbO,  in  which  the 
metal  is  divalent.*  Of  the  common  salts,  the  nitrate  and  acetate 
are  readily  soluble  in  water  ;  the  chloride,  bromide,  and  iodide  are 
sparingly  soluble. 

KHO,  NaHO,  or  NH4HO  gives  a  white  precipitate  of  either 
lead  hydroxide  or  a  basic  compound,  depending  upon  whether  the 
lead  solution  or  the  alkali  is  in  excess  all  the  time.  For  example, 
if  the  potash  be  added  to  the  lead  solution,  the  precipitation 

*  Salts  are  known  in  which  lead  is  tetravalent,  e.g.  lead  tetracetate. 
Pb(C2H3O2)4.  The  compound  PbCl4  has  also  been  obtained. 


So  Qualitative  Analysis. 

commences  in  the  presence  of  an  excess  of  lead  salt,  and  under  these 
conditions  a  basic  compound  is  formed.  If,  on  the  other  hand,  the 
lead  solution  is  added  to  the  alkali,  the  precipitation  is  made 
entirely  in  the  presence  of  excess  of  alkali,  and  then  the  hydroxide 
is  thrown  down,  Pb(HO)2.  The  precipitate  dissolves  in  excess  of 
KHO  or  NaHO,  but  not  in  NH4HO. 

K,CO3,  Na2CO3,  or  (NH4)2CO3  gives  a  precipitate  of  basic 
carbonate  of  lead. 

H2S  gives  a  black  precipitate  of  lead  sulphide,  PbS.  In  the 
presence  of  much  hydrochloric  acid,  the  precipitate  first  formed 
consists  of  a  brown  compound  having  the  composition  PbCl2,2PbS, 
which  by  the  further  action  of  H2S  is  converted  into  the  black  PbS. 

PbS  is  insoluble  in  cold  dilute  acids,  in  alkalies,  or  in  the  sul- 
phides of  the  alkalies.  It  is  readily  dissolved  by  hot  dilute  HNO3, 
giving  lead  nitrate  and  free  sulphur,*  the  latter  being  the  result 
of  a  secondary  reaction  between  the  liberated  H2S  and  a  further 
quantity  of  HNO3 ;  the  two  reactions  are  as  follows  : — 

PbS  +  2HNO3  =  Pb(NO3)2  +  H2S 
3H2S  +  2HN03  =  3S  +  4H20  +  2NO 

At  the  same  time  a  portion  of  the  PbS  is  oxidised  by  the  nitric 
acid  into  PbSO4. 

Strong  nitric   acid   converts  lead   sulphide   entirely   into   the 

sulphate. 

(NH4)2S  gives  the  same  precipitate. 

HoSO4  and  soluble  sulphates  give  a  white  precipitate  of  lead 
sulphate,  PbSO4.  Very  slightly  soluble  in  water ;  less  soluble  in 
the  presence  of  either  dilute  sulphuric  acid  or  alcohol ;  hence,  in 
very  dilute  solutions,  precipitation  is  accelerated  by  the  addition  of 
alcohol.  PbSO4  dissolves  by  long  boiling  with  strong  HC1,  yield- 
ing PbCL,.  It  dissolves  more  readily  in  strong  ammoniacal  solutions 
of  ammonium  acetate  or  tartrate,  as  well  as  in  hot  KHO  or  NaHO. 
From  these  it  is  again  precipitated  on  addition  of  H2SO4. 

HC1  and  soluble  chlorides  precipitate  white  lead  chloride, 
PbCl2 ;  the  precipitation  is  more  complete  in  the  presence  of  free 
nitric  or  hydrochloric  acids,  in  which  the  chloride  is  less  soluble 
than  in  water.  PbCl2  dissolves  moderately  freely  in  boiling  water 
(about  4  parts  in  100),  and  on  cooling,  the  solution  deposits  the 
compound  in  long  needle-shaped  crystals.  At  o°  the  liquid  holds 
in  solution  o'8  parts  in  100  of  water. 

*  The  sulphur  is  deposited  in  a  semi-pasty  condition,  and  often  appears 
dark-coloured,  even  black,  through  the  presence  of  particles  of  lead  sulphide 
which  are  entangled  in  it. 


Group  IL — Division  i.  81 

XI  gives  a  yellow  precipitate  of  PbI2.  Soluble,  but  to  a  less 
extent  than  the  chloride,  in  boiling  water  to  a  colourless  solution. 
Soluble  in  warm  acetic  acid,  from  which,  on  cooling,  it  crystallises 
in  golden  spangles. 

K,CrO4  precipitates  yellow  lead  chromate,  PbCrO4,  insoluble 
in  acetic  acid.  Soluble  in  dilute  HNO3  and  in  caustic  alkalies  (see. 
p.  42). 

Bismuth,  Bi. 

DRY  REACTIONS. — Bismuth  compounds  are  easily  reduced  when 
heated  with  Na2CO3  upon  charcoal.  The  metal,  however,  rapidly 
oxidises  when  strongly  heated,  hence  the  charcoal  becomes  covered 
with  an  incrustation  of  the  pale-yellow  oxide,  Bi2O3,  the  colour  of 
which  (as  is  the  case  with  most .  coloured  oxides)  appears  darker 
(orange-yellow)  while  hot.  Globules  of  the  metal,  if  detached  from 
the  charcoal,  may  be  at  once  distinguished  from  lead  or  silver  by 
their  brittleness.  Bismuth  dissolves  easily  in  HNO3,  but  is  scarcely 
attacked  by  HC1,  or  by  dilute  H2SO4. 

WET  REACTIONS. — Although  bismuth  forms  many  oxides,  only 
one  series  of  salts  is  known.  These  arc  derived  from  the  trioxide, 
Bi2O3.  All  the  other  oxides,  when  acted  upon  by  acids,  give  the 
same  series  of  salts,  in  which  the  bismuth  functions  as  a  trivalent 
element,  replacing  three  atoms  of  hydrogen.  In  the  case  of  the 
lower  oxide,  metallic  bismuth  is  deposited,  while  the  higher  oxides 
behave  as  peroxides,  evolving  oxygen  or  its  equivalent  in  chlorine  \ 
thus — 

3Bi202  +  6H2S04  =  2Bi2(S04)3  +  6H20 
Bi205  +  6HN03  -  02  +  2Bi(N03)3  +  3H2O 

Of  the  common  salts  of  bismuth  none  are  soluble  in  water  in 
the  ordinary  sense,  but  the  nitrate  and  chloride  are  readily  soluble 
in  water  acidified  with  the  respective  acids.  Water  alone,  converts 
these  salts  into  basic  compounds,  which  are  soluble  in  acid,  but  not 
in  water.  Thus,  when  water  is  added  to  bismuth  nitrate,  an  in- 
soluble basic  nitrate  is  formed,  and  nitric  acid  is  liberated.  This 
acid  is  capable  of  dissolving  the  basic  compound,  hence  the  reaction 
is  reversible,  and  is  therefore  incomplete  unless  the  nitric  acid 
generated  is  either  removed,  or  its  solvent  action  prevented  by 
considerable  dilution — 

Bi(NO3)3  +  2H2O  =  Bi(HO)2NO3*  +  2HNO3 

*  The  composition   of  this   compound  is   sometimes  represented  by   the 
formula  Bi(NO3)3,2Bi(HO)j.      It  may,   however,  be  more  simply  formulated 

G 


82  Qualitative  Analysis. 

In  the  case  of  bismuth  chloride,  the  oxychloride  is  thrown 
down — 

BiCl3  +  H2O  =  BiOCl  +  2HC1 

This  compound  is  not  so  easily  dissolved  by  HC1  as  the  basic 
nitrate  is  by  HNO3,  therefore  the  reaction  is  complete,  the  whole 
of  the  bismuth  being  precipitated  if  the  solution  is  dilute. 

KHO,  NaHO,  or  NH4HO  precipitates  the  white  hydroxide 
Bi(HO)3,  or  Bi2O3,3H2O.*  Insoluble  in  excess  of  the  precipitants. 
From  boiling  solutions,  or  on  heating  to  boiling,  the  monohydrated 
oxide  is  formed,  Bi2O3,H2O. 

K2CO3,  Na2CO3,  or  (NH4)2CO3  throws  down  a  white  basic 
carbonate,  (BiO)2CO3.t  Insoluble  in  excess  of  the  reagents. 

H2S  or  (NH4)2S  precipitates  bismuthous  sulphide,  Bi2S3,  as  a 
dark  brown,  almost  black,  compound.  Soluble  in  HNO3 ;  insoluble 
in  alkaline  sulphides.^ 

Sulphuric  acid  produces  no  precipitate  with  a  bismuth  salt. 
Potassium  dichromate,  K2Cr2O7,  precipitates  basic  bismuth  dichro- 
mate,  (BiO)2Cr2O7.  Insoluble  in  KHO.  (These  two  reactions 
distinguish  between  Pb  and  Bi.) 

The  most  characteristic  reaction  for  bismuth  is  the  formation 
of  the  insoluble  oxychloride,  BiOCl,  on  the  addition  of  water  to  an 
acid  solution  of  BiCl3.  As  explained  above,  the  reaction  with  the 
chloride  is  more  delicate  than  with  the  nitrate,  hence,  if  the  nitrate 
is  used,  it  should  be  converted  into  the  oxychloride.  This  can  be 
accomplished  by  the  addition  of  sodium  chloride,  whereby  both 
the  precipitated  basic  nitrate,  as  well  as  the  normal  nitrate  re- 
maining dissolved,  are  converted  by  double  decomposition  into 
oxychloride;  thus — 

Bi(NO3)3,2Bi(HO)3  +  3NaCl  =  sBiOCl  +  3H2O  +  3NaNO3 

or,  Bi(HO)2NO3  4-  NaCl  =  BiOCl  +  H2O  +  NaNO3 
Bi(NO3)3  +  3NaCl  +  H2O  =  BiOCl  +  3NaNO3  +  2HC1 

as  a  molecule  of  bismuth  nitrate,  in  which  two  of  the  NO3  groups  are  re- 
placed by  (HO) ;  thus,  Bi(HO)2NO3.  The  composition  in  both  cases  is  the 
same,  the  simpler  formula  being  merely  the  other  divided  by  three. 

*  Three  hydrated  oxides  are  known,  Bi2O3,3H2O,  Bi2O3,2H2O,  and 
Bi2O3,H2O. 

t  Or  this   may  be  regarded  as   normal    bismuth  carbonate  and    oxide, 

J  In  this  respect  bismuth  differs  from  the  elements  Sb  and  As,  with  which 
it  is  associated  in  the  "natural "  or  periodic  classification.  These  two  elements 
form  soluble  thioantimonates  and  thioarsenates,  no  such  compounds  of  bismuth 
being  known. 


Group  IL — Division  r.  83 

Cadmium,  Cd. 

DRY  REACTIONS. — Cadmium  compounds,  heated  on  charcoal 
with  sodium  carbonate,  are  easily  reduced  ;  but,  owing  to  the  ready 
volatility  of  the  metal,  the  latter  is  converted  into  the  oxide,  which 
is  deposited  as  a  brown  incrustation  upon  the  charcoal.* 

WET  REACTIONS. — The  metal  strongly  resembles  zinc  (with 
which  it  is  associated  in  the  natural  classification)  in  its  behaviour 
towards  acids,  dissolving  readily  in  dilute  acids  with  evolution  of 
hydrogen,  or,  in  the  case  of  nitric  acid,  of  oxides  of  nitrogen.  Only 
one  series  of  salts  is  known,  derived  from  the  only  oxide,  CdO.  Of 
the  common  salts,  the  nitrate,  sulphate,  chloride  (bromide,  iodide, 
and  acetate)  are  soluble  in  water. 

KHO,  NaHO,  or  NH4HO  precipitates  the  white  hydroxide, 
Cd(HO)2.  Insoluble  in  excess  of  KHO  or  NaHO,  but  soluble  in 
NH4HO. 

K,COv,  Na2CO3,  or  (NH4)2CO3  gives  a  white  precipitate  of 
CdCO3.  The  presence  of  NH4HO  prevents  the  precipitation.  The 
precipitate  is  not  soluble  in  excess  of  the  reagent  as  ordinarily 
used  ;  but  it  dissolves  in  concentrated  solutions  of  K2CO3  and 
Na2CO3.  From  this  solution  it  is  reprecipitated  on  dilution. 

HLS  or  (NH4)2S  precipitates  cadmium  sulphide,  CdS,  distin- 
guished from  the  sulphides  of  all  the  other  metals  of  this  division 
by  its  pure  yellow  colour.  It  is  more  easily  soluble  in  acids  than 
the  other  sulphides  of  the  group,  and  therefore  the  acid  solution 
from  which  it  is  precipitated  must  be  dilute,  and  not  too  strongly 
acid,  to  ensure  complete  precipitation. 

CdS  is  insoluble  in  potassium  cyanide,  therefore  it  is  capable  of 
being  precipitated  in  the  presence  of  this  salt.  And  for  the  same 
reason,  H2S  will  throw  down  a  precipitate  of  CdS  from  a  solution 
of  CdCy2  in  excess  of  KCy,  which  contains  the  double  cyanide 
CdCy2,2KCy  (see  Method  of  Separation  from  Copper).  CdS  is 
insoluble  in  alkaline  sulphides,  which  distinguishes  it  from  arsenious 
sulphide,  which  is  the  only  other  yellow  sulphide  (see  footnote, 
P-  77). 

Copper,  Cu. 

DRY  REACTIONS, — Copper  compounds  are  reduced  to  metallic 
copper  when  strongly  heated  upon  charcoal  along  with  sodium 

*  With  the  exception  of  the  rare  mineral  greenockite,  CdS,  cadmium  is 
always  found  in  nature  closely  associated  with  zinc  ores,  and  only  in  small 
quantities.  Being  much  more  volatile,  however,  than  zinc,  it  is  often  possible 
to  obtain  the  brown  incrustation  of  CdO  before  enough  of  the  zinc  has  been 
vaporised  to  mask  the  reaction. 


84  Qualitative  Analysis. 

carbonate  in  the  reducing  flame.  Reddish  scales,  or  even  globules, 
of  metal  will  be  found.  Heated  in  a  borax  bead,  copper  salts 
impart  a  colour  which  is  green  while  the  bead  is  hot,  but  bluish 
when  cold.  A  more  delicate  dry  test  for  copper  is  made  by  heating 
the  compound  upon  a  platinum  loop  in  a  Bunsen  flame,  ana; 
supplying  the  flame  at  the  same  time  with  a  little  hydrochloric  acid 
gas,  which  is  admitted  by  one  of  the  air-holes  of  the  lamp.  A  little 
strong  hydrochloric  acid  is  heated  in  a  test-tube  having  a  delivery 
tube  leading  into  one  of  the  air-holes,  in  the  manner  shown  in 
Fig.  n.  Under  these  circumstances,  the  copper  compound,  which 


FIG.  ii. 

may  otherwise  impart  no  colour  to  the  flame,  gives  a  brilliant  blue 
colour,  which  instantly  changes  to  green  when  the  supply  of  acid 
gas  is  momentarily  stopped  by  partially  withdrawing  the  test-tube. 
WET  REACTIONS. — Copper  is  not  acted  upon  by  dilute  HC1  or 
dilute  H2SO4.  Boiling  strong  HC1  slowly  dissolves  it,  giving 
cuprous  chloride  and  hydrogen.  Hot  concentrated  H2SO4  gives 
SO2  and  copper  sulphate.  Nitric  acid  readily  dissolves  the  metal, 
forming  Cu(NO3)2  and  oxides  of  nitrogen.  Copper  forms  two 
series  of  salts,  cuprous  and  cupric,  derived  from  the  two  oxides 
Cu2O  and  CuO.  The  former  readily  pass  by  oxidation  into  cupric 
compounds. 


Group  II. — Division  I.  85 


(a}  Cupric  Salts.  —  Of  the  common  cupric  salts,  the  sulphate, 
nitrate,  chloride  (bromide  and  acetate)  are  readily  soluble  in  water. 
In  the  crystallised  or  hydrated  condition  they  are  blue  or  green, 
but  in  the  anhydrous  state  either  white  or  pale  yellow. 

KHO  or  NaHO  produces  a  pale  blue  precipitate  of  cupric 
hydroxide  or  hydrated  oxide  Cu(HO)2  or  CuO,H2O.  Insoluble  in 
excess.*  On  warming  the  mixture  (or  if  the  precipitation  is  con- 
ducted with  hot  solutions),  this  loses  a  portion  of  its  water  of 
hydration  and  is  changed  to  a  nearly  black  hydrated  oxide  — 

3CuO,H2O  or  Cu(HO)2,2CuO 

NH,HO  or  (NH4)2CO3  precipitates  a  light  blue  basic  com- 
pound, readily  soluble  in  excess  to  a  deep  blue  solution  (charac- 
teristic of  copper  compounds).  If  the  blue  solution  be  allowed  to 
evaporate,  it  deposits  dark  blue  crystals,  having  the  composition 
CuSO4,4NH3,H2O  f  in  the  case  of  the  sulphate;  and  CuCl2,4NH3,H2O 
with  the  chloride. 

K  CO  or  Na2CO3  gives  a  greenish  precipitate  of  the  basic 
carbonate,  Cu(CO3),Cu(HO)24  Insoluble  in  excess  of  the  reagent 
under  ordinary  circumstances.  If,  however,  the  reagent  be  a  highly 
concentrated  solution  of  the  alkaline  carbonate,  the  precipitate  dis- 
solves to  a  deep  blue  solution,  believed  to  contain  a  double  carbonate 
of  copper  and  the  alkali.  Dilution  with  water  decomposes  this 
compound,  and  reprecipitates  the  basic  carbonate  (see  also  Cd,  Ni, 
Co,  Fe).  On  boiling  the  mixture,  the  precipitate  is  converted  into 
the  black  hydrated  oxide  ;  thus  — 

3[CuC03)Cu(HO)2]  -  2[Cu(HO)2,2CuO]  +  3CO2  +  H2O 


H2S  or  (NH4)2S  produces  a  nearly  black  precipitate  of  cupric 
sulphide,  CuS,§  which,  when  exposed  to  the  air  in  a  moist  condition, 
absorbs  oxygen  and  is  converted  into  the  sulphate.  The  precipitate 
is  slightly  soluble  in  ammonium  sulphide.  It  readily  dissolves  in 
potassium  cyanide,  therefore  H2S  fails  to  give  a  precipitate  of 
copper  sulphide  from  a  solution  of  Cu2Cy2  1|  in  KCy  (which 
contains  cuprous  potassium  cyanide,  Cu2Cy2,6KCy)  (compare 
cadmium). 

*  In  the  presence  of  tartaric  acid  (or  alkaline  tartrates),  Cu(HO)2  is  dissolved 
by  excess  of  KHO  (Feeling's  solution.  See  below). 

f  Copper  sulphate  crystals  themselves  have  the  composition  CuSO4,5H2O. 

J  This  compound  occurs  in  nature  as  the  mineral  malachite. 

§  The  composition  of  precipitated  cupric  sulphide  is  said  to  be  represented 
by  the  formula  Cu4S3  (Thomson). 

||  See  cuprous  cyanide. 


86  Qualitative  Analysis. 

Reduction  of  Cupric  Salts.— (i)  Many  organic  substances 
reduce  cupric  salts  in  alkaline  solutions,  with  precipitation  of 
cuprous  oxide,  Cu2O.  Thus,  if  KHO  be  added  to  a  solution  of 
CuSO4  in  the  presence  of  grape  sugar,  the  Cu(HO)2  first  precipi- 
tated dissolves  in  excess  of  KHO,  giving  a  blue  solution  ;  on  gently 
warming  the  liquid,  a  bright  red  precipitate  of  Cu2O  is  thrown 
down.*  The  reaction  may  be  regarded  as  an  abstraction  of  one 
atom  of  oxygen,  by  the  sugar,  from  two  molecules  of  cupric  oxide — 
2CuO  =  Cu2O  +  O 

Cuprous  oxide  is  soluble  in  hydrochloric  acid,  with  the  formation 
of  cuprous  chloride. 

(2)  Cupric  salts  may  be  reduced  by  nascent  hydrogen.  Thus, 
when  metallic  copper  is  placed  in  a  solution  of  cupric  chloride  in 
strong  hydrochloric  acid,  and  the  mixture  boiled,  the  green  cupric 
chloride  becomes  colourless,  owing  to  its  conversion  into  cuprous 
chloride  by  the  hydrogen  disengaged  by  the  action  of  the  HC1 
upon  the  copper — 

2CuCl2  +  H2  =  Cu2Cl2  f  +  2HC1 

If  the  colourless  solution  be  poured  into  water,  the  cuprous  chloride 
is  thrown  down  as  a  white  precipitate  ;  insoluble  in  water,  soluble 
in  HC1,  in  NH4HO,  and  in  NH4C1. 

(ft)  Cuprous  Salts. — The  common  cuprous  salts  are  all 
insoluble  in  water.  P'or  the  reactions,  a  solution  of  cuprous 
chloride  in  hydrochloric  acid  may  be  used. 

KHO  or  NaHO  gives  a  yellow  precipitate  of  cuprous  hydroxide, 
Cu2(HO)2  or  Cu2O,H2O.  If  the  mixture  be  heated,  the  precipitate 
is  converted  into  the  red  cuprous  oxide. 

NH.HO  gives  no  precipitate,  but  forms  a  soluble  compound 
having  the  composition  Cu2Cl2,2NH8.  The  solution  is  colourless, 
but  has  the  property  (also  possessed  by  the  solution  of  Cu2Cl2  in 
HC1)  of  absorbing  oxygen  from  the  air,  first  becoming  brown,  and 
finally  depositing  a  greenish  precipitate  of  cupric  oxychloride, 
CuCl2,3CuO,4H204 

*  This  reaction  is  utilised  as  a  test  for  sugar.  The  reaction  is  more  delicate 
when  an  alkaline  solution  of  cupric  tartrate  (Feeling's  solution)  is  employed. 
This  is  prepared  by  neutralising  a  solution  of  tartaric  acid  with  excess  of 
KHO,  and  adding  to  the  alkaline  liquid  a  small  quantity  of  copper  sulphate 
solution.  A  light  blue  solution  is  obtained. 

t  The  cuprous  and  cupric  compounds  are  related  to  each  other  in  the  same 
way  as  the  mercurous  and  mercuric  salts.  Some  chemists  regard  the  cuprous 
salts  as  containing  monovalent  copper.  It  is  preferable,  however,  to  represent 
them  as  compounds  of  the  divalent  double  atom  Cu2,  yielding  on  solution 
monovalent  ions  Cu.  Cu. 

J  This  ammoniacal  solution  of  cuprous  chloride  also  absorbs  carbon  mon- 


Separation  of  the  Metals  of  Group  II. — Division  i .     87 

KI  produces  a  yellowish-white  precipitate  of  cuprous  iodide, 
Cu2I2- 

Cu2Cl2  +  2KI  =  Cu2I2  +  2KC1 

Cupric  iodide  is  unknown.  When  KI  is  added  to  a  solution  of  a 
cupric  salt,  cuprous  iodide  is  thrown  down,  and  iodine  is  liberated, 
which  colours  the  liquid  brown — 

2CuCl2  +  4KI  =  Cu2I2  +  4KC1  +  I2 

In  the  presence  of  a  reducing  agent,  such  as  a  ferrous  salt,  or 
sulphurous  acid,  the  cupric  salt  is  first  reduced  to  the  cuprous 
state,  which  with  the  KI  then  gives  cuprous  iodide  without  separa- 
tion of  iodine — 

2CuCl2  +  2FeCl2  -  2FeCl3  +  Cu2Cl2 
2CuCl2  +  H2SO3  +  H2O  =  H2SO4  +  2HC1  +  Cu2Cl2 

KCy  gives  a  white  precipitate  of  Cu3Cy2 ;  soluble  in  excess  of 
KCy,  giving  a  double  cyanide,  Cu2Cy2,  6 KCy.* 

[Cupric  cyanide,  although  known,  is  very  unstable,  hence  when 
J£Cy  is  added  to  a  cupric  salt,  the  cupric  cyanide,  even  if  formed, 
quickly  decomposes,  and  a  mixture  or  compound  of  cuprous  cyanide 
and  cupric  cyanide  is  produced.  As  with  the  iodide,  if  reducing 
agents  are  present,  cuprous  cyanide  alone  is  produced.] 

SEPARATION  OF  THE  METALS  OF  GROUP  II.— DIVISION  i. 

The  separation  of  these  metals  is  based  upon — 

1.  The  difference  in  the  behaviour  of  their  sulphides  towards 
nitric  acid  ;   sulphides  of  Cu,  Cd,  and  Bi  being  dissolved,  while 
those  of  Hg  and  Pb  are  either  not  dissolved  or  are  changed  into 
other  insoluble  compounds. 

2.  The  solubility  of  the  hydroxides  of  Cd  and  Cu  in  NH4HO, 
and  the  insolubility  of  Bi(HO)3. 

3.  The  solubility  of  copper  sulphide  in  KCy,  and  the  insolubility 
of  CdS. 

The  solution  containing  salts  of  these  metals  is  acidified  by  the 
addition  of  a  few  drops  of  dilute  HCl,t  and  a  stream  of  H2S  is 

oxide,  forming  a  compound  believed  to  have  the  composition  COCu2Cl2(2H2O. 
It  likewise  absorbs  acetylene,  C2H2,  producing  a  red  precipitate  of  cuprous 
acetylide  ;  thus — 

Cu2Cl2,2NH3  +  H2O  +  C2H2  =  Cu2C2,H2O  +  2NH4C1 

*  In  this  solution  the  copper  exists  in  the  cyanogen  anion,  probably  Cu^Cy3. 

t  In  the  ordinary  course  of  analysis,  the  metals  of  Group  I.  are  thereby 
precipitated  as  chlorides ;  nearly  the  whole  of  the  lead  will  therefore  be 
removed  before  Group  II.  is  examined.  When  the  exercise  is  confined  to  this 
group,  any  precipitate  of  PbCl2  obtained  may  be  removed  by  filtration  and 
examined  separately. 


88 


Qualitative  Analysis. 


.allowed  to  bubble  moderately  slowly  through  the  solution  until 
precipitation  is  complete  (see  Precipitation,  p.  7).  During  the 
process  the  liquid  should  be  frequently  stirred  with  the  tube 
delivering  the  gas.  After  filtration,  the  solution  should  be  diluted 
with  water,  and  H2S  should  be  again  passed  through  the  liquid  for 
a  moment  or  two,  to  make  sure  that  the  precipitation  has  been 
complete, 


The  precipitate  contains  the  sulphides  of  all  the  metals  of  the  division.  It 
shoald  be  well  washed,*  and  then  transferred  to  a  small  porcelain  dish 
with  the  least  ^possible  quantity  of  water.  About  an  equal  volume  of 
strong  HNO3  is  then  added,  and  the  mixture  boiled  until  no  further 
dissolving  action  can  be  detected.  The  mixture  is  then  diluted,  and  a 
few  drops  of  dilute  H2SO4  added.  Before  filtering,  the  mixture  should 
be  cooled. 


The  residue  may  contain  HgS 
(black),  Hg(NO3)2,2HgS  (white), 
and  PbSO4  (white).  Boil  with  a 
solution  of  ammonium  acetate. 


The  filtrate  contains  the  nitrates 
of  Bi,  Cd,  and  Cu.  Add  excess 
ofNH4HO,  and  boil.  \ 


The      precipi- 

The solution  con- 

The   residue. 

The  solution. 

tate    is  white 

tains     ammonia- 

Dissolve   in  the 

Add  K2CrO4. 

Bi2O3,H2O.t 

cal    Cd   and   Cu 

least       quantity 

A  yellow  pre- 



compounds.  The 

of   aqua     regia. 

cipitate        of 

Confirm  by  dis- 

presence of  Cu  is 

Boil     to    expel 

PbCrO4. 

solving  in  a  few 

seen  by  the  blue 

chlorine,       and 

drops  of  HC1, 

colour.     If  blue, 

neutralise     with 

and  adding  the 

add    KCy    until 

NaHO.J  Acidify 

solution     drop 

colourless,      and 

with  I1C1,   and 
introduce  a  strip 
of  clean  copper, 

by  drop  into  a 
test-tube  nearly 
filled          with 

passH2S.  Yellow 
CdS  precipitated. 

which   will    be- 
come        coated 

water.     White 
BiOCl  is  pre- 

If copper  is  absent, 
H2S       may      be 

with    a    silvery 

cipitated. 

added  at  once. 

deposit  of  mer- 

cury. 

*  Unless  HC1  (and  soluble  chlorides)  be  washed  out  of  this  precipitate,  the 
addition  of  HNO3  will  result  in  the  formation  of  a  little  aqua  regia,  and  this 
will  dissolve  a  portion  of  the  mercuric  sulphide. 

t  If  any  lead  sulphate  escaped  precipitation  in  the  previous  step,  or  if  any 
of  the  mercuric  sulphide  was  dissolved  (see  previous  note),  compounds  of  these 
metals  will  be  here  thrown  down  by  NH4HO.  Hence  it  is  necessary  to  con- 
firm Bi  as  indicated. 

%  This  step  is  necessary,  in  order  to  remove  HNO3,  the  presence  of  which 
might  prevent  the  deposition  of  Hg. 


CHAPTER  IX. 

REACTIONS  OF  THE  METALS  OF  GROUP  II. — DIVISION  2. 

Arsenic,  Antimony,  Tin  [Gold,  Platinum]. 

THE  two  elements  arsenic  and  antimony  belong  to  the  same 
natural  family.  Tin,  on  the  other  hand,  is  more  nearly  related  to 
lead.  The  three  elements  are  associated  together  in  the  same 
analytical  group,  for  the  reason  that  they  possess  in  common  the 
property  of  forming  "  thio  "  acids,  whose  alkaline  salts  are  soluble  in 
water  ;  namely,  thio-arsenites,  thio-antimonites  and  thio-stannates. 
In  other  words,  the  sulphides  of  these  metals  are  soluble  in  alkaline 
sulphides. 

Arsenic,  As. 

DRY  REACTIONS. — Compounds  of  arsenic  are  easily  reduced 
and  the  element  obtained  in  the  "  metallic  "  state,  by  heating  them 
with  suitable  reducing  agents. 

Thus,  when  heated  upon  charcoal  with  Na2CO3  and  KCy, 
arsenical  compounds  are  reduced ;  but  the  metal,  being  extremely 
volatile  and  readily  combustible,  is  for  the  most  part  burnt  to 
arsenious  oxide,  As4O6,  which  passes  off  as  a  white  fume.  At  the 
same  time  some  of  the  vapour  of  the  element  itself  is  carried  away 
with  the  fumes  of  the  oxide,  and  is  readily  recognised  by  its  strong, 
unpleasant,  and  characteristic  garlic-like  odour.  Arsenic  cannot  be 
melted  by  heat  after  the  manner  of  most  metals,  but  passes  from 
the  solid  to  the  vaporous  states  without  liquefying ;  hence  it  never 
yields  metallic  globules  upon  the  charcoal. 

The  reduction  may  be  made  by  heating  the  arsenic  compound 
in  a  glass  tube  with  KCy,  or  a  mixture  of  Na2CO3  and  KCy.  The 
reaction  is  conveniently  studied  by  using  arsenious  oxide.  A  small 
fragment  (about  the  size  of  a  pin's  head)  is  placed  in  a  narrow  test- 
tube  *  and  covered  by  adding  a  mixture  of  Na2CO3  and  KCy  (equal 

*  For  such  experiments  small  test-tubes  4  X  tV  inches  answer  admirably. 
Bulb  tubes  are  neither  necessary  nor  desirable. 


go  Qualitative  Analysis. 

parts),  the  materials  being  as  dry  as  possible.  The  total  quantity 
of  material  in  the  tube  should  not  occupy  more  space  than  is  shown 
in  Fig.  12.  On  the  application  of  a  gentle  heat,  the  first  effect  will 
be  the  expulsion  of  moisture  from  the  imperfectly  dried  materials, 


FIG.  12. 

which  condenses  upon  the  sides  of  the  tube.  This  may  be  driven 
up  the  tube  by  gently  warming  it,  and  finally  removed  by  intro- 
ducing a  "spill"  of  blotting-paper.  When  no  more  moisture 
collects,  the  mixture  may  be  steadily  heated  in  the  tip  of  a  small 
Bunsen  flame.  The  arsenic  under  these  circumstances  is  vaporised 
without  undergoing  combustion,  and  sublimes  upon  the  tube  as  a 
metallic  mirror.  Sufficient  of  the  vapour  escapes  condensation  to 
enable  the  strong  garlic  odour  to  be  detected. 

The  reaction  which  takes  place  may  be  expressed  by  the 
equation — 

As4O6  +  6KCy  =  As4  +  6KOCy 

The  oxide  of  arsenic  may  be  reduced  by  being  heated  with  charcoal 
alone  in  a  glass  tube,  when  a  metallic  sublimate  is  also  obtained. 
In  this  case  it  is  necessary  to  cover  the  arsenic  compound  with  a 
layer  of  charcoal,  which  should  be  strongly  heated  before  the  oxide 
of  arsenic  becomes  hot,  so  that  the  vapour  shall  pass  through  the 
heated  carbon  ;  otherwise  the  arsenious  oxide  might  be  entirely 
volatilised  before  the  carbon  had  become  hot  enough  to  reduce  the 
compound. 

WET  REACTIONS.— The  element  arsenic  has  little  or  no  claim  to 
be  ranked  as  a  metal  in  the  strict  acceptation  of  the  term.  Being 
like  a  metal  in  appearance  and  in  some  of  its  physical  properties, 
it  is  usually  called  a  metalloid.  In  the  natural  family  to  which  it 
belongs,  it  stands  between  phosphorus  and  nitrogen  on  the  one 


Group  II.— Division  2.  91 

hand,  and  antimony  and  bismuth  on  the  other,  showing  in  its 
chemical  habits  a  strong  similarity  to  phosphorus. 

It  forms  two  oxides,  both  of  which  are  acidic  in  their  character  ; 
and  all  the  salts  of  arsenic  are  such  as  contain  this  element  in  the 
acidic  or  negative  portion  of  the  molecule,  united  with  other  metals 
as  the  base  ;  such,  for  example,  as  the  arsenites  and  arsenates  of 
various  metals.  No  oxysalts  of  arsenic  are  known  in  which  the 
element  plays  the  part  of  a  base,  such  as  nitrate,  sulphate,  carbonate, 
etc.,  of  arsenic. 

(a)  Arsenious  compounds,  derived  from  arsenious  oxide, 
As4O6.*  The  arsenites  of  sodium,  potassium,  and  ammonium  alone 
are  soluble  in  water.  For  the  following  reactions,  potassium  arsenite, 
K3AsO3,  or  a  solution  of  As4O6  in  dilute  HC1,  may  be  employed. 

H2S  or  (NH4)2S  precipitates  from  slightly  acid  solutions  yellow 
arsenious  sulphide,  As2S3,  soluble  in  excess  of  ammonium  sulphide, 
giving  ammonium  thio-arsenite  ;  thus — 

As2S3  +  3(NH4)2S  -  2(NH4)3AsS3 

With  yellow  ammonium  sulphide  (polysulphide)  the  arsenious 
sulphide  is  dissolved,  with  the  formation  of  ammonium  thio-arsenate, 
the  As2S3  first  uniting  with  the  sulphur  of  the  polysulphide  to  form 
As2S5.  From  this  solution  the  higher  sulphide  is  precipitated  on 
the  addition  of  an  acid.  If  H2S  be  passed  through  an  aqueous 
solution  of  As4O6  no  precipitate  is  produced,  but  the  liquid  becomes 
yellow  owing  to  the  presence  of  arsenious  sulphide  in  solution  in 
the  colloidal  state.  The  addition  of  HC1  causes  the  precipitation 
of  the  yellow  sulphide.  From  neutral  or  alkaline  solutions,  H2S 
(or  (NH4)2S)  gives  no  precipitate,  as  under  these  circumstances  the 
soluble  thio  salt  is  produced.  The  addition  of  an  acid  decomposes 
the  thio  salt,  and  arsenious  sulphide  is  thrown  down.  The  reactions 
may  be  expressed  thus — 

K3As03  +  3H2S  =  K3AsS3  +  3H2O 
2K3AsS3  +  6HC1  =  6KC1  +  3H2S  +  As2S3f 

Arsenious  sulphide  is  soluble  in  caustic  alkalies  and  ammonia,  also 

*  The  acid  corresponding  to  this  oxide  is  not  known.  Three  series  of  salts 
are  known,  which  may  be  considered  as  being  derived  from  the  three  hypo- 
thetical acids — ortho-arsenious  acid,  H3AsO3  ;  pyro-arsenious  acid,  H4As2O5  ; 
and  metarsenious  acid,  HAsO2. 

t  The  hypothetical  thio-arsenious  acid,  H3AsS3,  may  be  supposed  to  be 
first  formed,  and  to  at  once  split  up  into  the  thio-anhydride  As2S3  and  into 
H2S,  just  as  on  the  addition  of  an  acid  to  a  carbonate  the  unstable  carbonic 
acid  H2CO3  breaks  up  into  the  anhydride  COj  and  water. 


92  Qualitative  Analysis. 

in  ammonium  carbonate,  forming  in  each  case  a  mixture  of  arsenite 
and  thio-arsenite  of  the  alkali  ;  thus— 

As2S3  +  6KHO  =  K3AsO3  +  K3AsS3  +  3H2O 
As2S3  +  3(NH4)2C03  =  (NH4)3As03  +  (NH4)3AsS3  +  3CO2 

On  the  addition  of  acid,  the  mixed  •  arsenite  and  thio-arsenite  are 
decomposed,  and  arsenious  sulphide  is  reprecipitated — 

K3AsO3  +  K3AsS3  +  6HC1  =  6KC1  +  3H2O  +  As2S3 

Arsenious  sulphide  is  practically  insoluble  in  HC1  (contrast  Sb2S3), 
but  readily  dissolves  in  HNO3,  or  in  HC1  with  addition  of  a  crystal 
of  KC1O3  ;  being  oxidised  in  each  case  into  arsenic  acid.* 

CuSO4  produces,  in  a  solution  of  potassium  arsenite,  a  green 
precipitate  of  hydrogen  cupric  arsenite,  HCuAsO3;f  soluble  in 
ammonia  and  caustic  alkalies.  If  the  solution  be  boiled,  the  arsenite 
is  oxidised  to  arsenate  at  the  expense  of  the  copper,  which  is  thereby 
reduced  to  cuprous  oxide,  and  precipitated  in  this  form  ;  thus — 

2CuSO4  +  K3AsO3  +4KHO  =  K3AsO4  +  2K2SO4  +  2H2O  +  Cu2O 

AgNO ;  precipitates,  from  a  solution  of  potassium  arsenite,  pale- 
yellow  silver  arsenite,  Ag3AsO3,  soluble  both  in  NH4HO  and  in 
HNO3.  If  the  precipitate  be  dissolved  in  either  of  these  solvents, 
it  can  only  be  reprecipitated  by  neutralisation  with  the  other,  when 
the  utmost  care  has  been  taken  to  avoid  excess  of  the  first,  because 
the  silver  arsenite  is  soluble  also  in  ammonium  nitrate.  When  the 
ammoniacal  solution  is  boiled  for  some  time,  metallic  silver  is 
precipitated,  and  the  arsenite  is  oxidised  to  arsenate.  The  solution 
of  silver  arsenite  in  ammonia,  and  the  decomposition  on  boiling, 
may  be  expressed  by  the  following  equations  : — 

Ag3AsO3  +  2NH4HO  =  (NH4)2AgAsO3  +  Ag2O  $  +  H2O 
(NH4)2AgAs03  +  Ag20  -  (NH4)2AgAs04  +  Ag2 

Precipitation  by  Copper  (Reinsch's  test).— If  a  strip  of  clean 
copper  foil  be  introduced  into  a  solution  of  arsenious  oxide  in  HC1, 
or  an  arsenite  acidified  with  the  same  acid,  and  the  mixture  be 

*  The  modus  operandi  of  this  oxidation  of  the  sulphide  by  HC1  and  KC1O3 
is  that  the  chlorine  first  converts  the  sulphide  into  arsenic  trichloride,  which,  in 
presence  of  water,  forms  arsenious  acid,  and  that  the  chlorine  oxidises  this  into 
arsenic  acid.  The  changes  may  be  thus  expressed — 

(i)  As2S3  +  3C12  =  2AsCl3  +  38 

(a)  2AsCl3  +  6H2O  =  6HC1  +  2H3As03 

(3)  H3AsO3  +  H2O  +  C12  =  2HC1  +  H3AsO4 

f  Used  as  a  pigment,  under  the  name  of  Scheele  s  green. 
%  Ag2O  dissolves  in  ammonia. 


Group  II. — Division  2.  93 

warmed,  metallic  arsenic  is  deposited  upon  the  copper,  at  the  same 
time  uniting  with  it,  forming  copper  arsenide,  Cu5As2.  If  the 
copper  be  then  dried,  and  gently  heated  in  a  dry  test-tube,  the 
arsenic  will  be  volatilised,  and  at  the  same  time  oxidised,  giving, 
therefore,  a  white  crystalline  sublimate  of  As4O6  (contrast  antimony]. 

(b)  Arsen/c  compounds,  derived  from  arsenic  pentoxide, 
AsaO5.*  Arsenates  of  sodium,  potassium,  and  ammonium  are 
soluble  in  water. 

H2S. — From  neutral  or  alkaline  solutions  of  an  arsenate  no 
precipitate  is  produced,  the  soluble  thio-arsenate  being  formed. 
The  addition  of  HC1  to  this  solution  throws  down  the  pentasulphide 
in  the  form  of  a  yellow  precipitate — 

K3As04  +  4H2S  =  4H20  +  K3AsS4 
2K3AsS4  +  6HC1  =  6KC1  +  As2S5  +  sH2S 

From  acidified  solutions  of  an  arsenate,  H2S  gives  a  precipitate 
after  a  short  time,  which  is  either  As2S5  or  a  mixture  of  As2S3  and 
S,  depending  upon  conditions. 

If  the  solution  is  strongly  acid,  and  the  gas  is  passed  rapidly, 
the  precipitate  which  slowly  comes  down  is  the  pentasulphide. 
On  the  other  hand,  if  the  solution  is  less  strongly  acid,  and  the 
H2S  is  passed  slowly,  the  arsenic  acid  is  first  reduced  to  arsenious 
acid,  with  deposition  of  sulphur,  and  the  arsenious  acid  as  it  forms 
is  converted  into  arsenious  sulphide  ;  thus — 

(1)  H3AsO4  +  H2S  =  H2O  +  S  +  H3AsO3 

(2)  2H3AsO3  +  3H2S  =  As2S3+  6H2O 

This  reducing  action  of  H2S  is  very  slow,  therefore  complete  pre- 
cipitation requires  considerable  time.  Warming  the  liquid  hastens 
the  action.  The  addition  of  a  more  powerful  reducing  agent,  such 
as  sulphurous  acid,  produces  the  effect  at  once. 

As2S5  dissolves  in  alkaline  sulphides,  forming  thio-arsenates, 
similar  to  the  thio-arsenites. 

CuSO,  gives  a  pale  bluish  precipitate  of  hydrogen  cupric 
arsenate,  HCuAsO4.  Soluble,  like  the  corresponding  arsenite,  in 
ammonia  ;  but  the  copper  is  not  reduced  on  heating  the  solution, 
for  the  reason  that  arsenates  are  incapable  of  any  further  oxidation 
— in  other  words,  they  do  not  act  as  reducing  agents  (contrast 
arsenites). 

*  Three  arsemV  acids  derived  from  this  oxide  are  known,  namely,  ortho- 
arsenic  acid,  H3AsO4 ;  pyro-arsenic  acid,  HsAseO?  ;  and  metarsenic  acid, 
HAsO3.  The  two  latter,  when  dissolved  in  water,  are  converted  into  the  ortho- 
acid  ;  it  is,  therefore,  only  possible  to  have  an  aqueous  solution  of  the  ortho-acid. 


94  Qualitative  Analysis. 

AgNO,  produces  a  chocolate-coloured  precipitate  of  silver 
arsenate,  Ag3AsO4,  which,  like  the  arsenite,  dissolves  in  NH4HO, 
in  HNO3,  and  in  NH4NO3.  If  a  mixture  of  silver  arsenate  and 
arsenite  be  carefully  dissolved  in  HNO3,  avoiding  any  excess,  and 
then  ammonia  added  drop  by  drop,  the  silver  arsenate  is  pre- 
cipitated first  (recognised  by  its  chocolate  colour),  and  afterwards 
the  yellow  arsenite. 

The  ammoniacal  solution  of  silver  arsenate  is  not  reduced  on 
boiling,  for  the  same  reason  that  the  copper  salt  is  not  reduced 
(contrast  arsenites]. 

MgSO4,  in  presence  of  NH4C1  and  NH4HO,  gives  a  white 
crystalline  precipitate  of  ammonium  magnesium  arsenate, 
(NH4)MgAsO4  *  ;  practically  insoluble  in  water.  (The  correspond- 
ing arsenite  is  known,  but,  being  readily  soluble  in  water,  it  is  not 
produced  by  precipitation  ;  hence  this  reaction  serves  to  distinguish 
an  arsenate  from  an  arsenite?) 

Marsh's  Test. 

In  the  presence  of  nascent  hydrogen,  both  arsenic  and  arsenious 
compounds  are  reduced,  and  arsenuretted  hydrogen  is  evolved.  Thus, 
if  a  solution  of  arsenious  or  arsenic  oxide  be  subjected  to  electrolysis, 
or  if  such  solutions  are  introduced  into  a  mixture  from  which  hydro- 
gen is  being  generated  (e.g.  zinc  or  magnesum  with  dilute  acid),  this 
compound  of  arsenic  and  hydrogen  is  produced.  The  action  may 
be  regarded  as  taking  place  in  two  stages,  first  the  reduction  of  the 
arsenical  compound  to  metallic  arsenic,  and  then  the  further  action 
of  the  nascent  hydrogen  upon  this  ;  thus — 

As4O6  +  6H2  =  As4  +  6H2O 
As4  +  6H2  =  4AsH8 

The  properties  of  arsenuretted  hydrogen  which  are  made  use  of  in 
analysis  are  the  following  : — 

(1)  The  deposition  of  metallic  arsenic  f  from  the  flame  of  the 
burning  gas  when  a  cold  object  is  depressed  upon  the  flame. 

(2)  The  decomposition  of  the  compound  on  passing  through  a 
heated  tube,  with  deposition  of  an  arsenical  mirror. 

*  This  compound  closely  resembles  the  corresponding  phosphate, 
(NH4)MgPO4,  and  is  precipitated  under  the  same  conditions.  It  may,  how- 
ever, be  at  once  distinguished  from  the  phosphate  by  dissolving  it  in  HC1,  and 
adding  H2S,  when  a  yellow  precipitate  of  As2S5  is  produced. 

f  Recent  experiments  seem  to  prove  that  this  deposit  is  in  reality  a  solid 
hydride  of  arsenic  AsH  (Retgers,  ZfitS.f,  anorg.  Chem.,  iv.  739). 


Group  II. — Division  2. 


95 


(3)  The  action  of  the   gas   upon   a   solution  of  silver  nitrate, 
resulting  in  the  precipitation  of  metallic  silver. 

The  reaction  is  made  in  a  small  hydrogen  generating  apparatus, 
preferably  a  Woulff's  bottle,  of  about  200  cub.  cms.  capacity.  In  this 
hydrogen  is  slowly  gene- 
rated from  zinc  and  dilute 
sulphuric  acid,  both  mate- 
rials being  free  from  arse- 
nic. To  the  exit-tube  is 
attached  a  tube,  drawn 
out  of  a  piece  of  combus- 
tion tube  in  such  a  manner 
as  to  present  one  or  two 
constricted  places  in  its 
length,  as  shown  in  Fig.  13. 
As  soon  as  the  air  is  all 
expelled  from  the  appa- 


ratus, the  issuing  hydrogen 
is  inflamed.* 

A  small  quantity  of  the 
arsenical  solution  is  now 
introduced  through  the 

thistle-tube.  The  first  effect  of  this  is  to  cause  the  precipitation 
of  arsenic  upon  the  zinc  ;  which,  constituting  a  voltaic  couple,  at 
once  gives  rise  to  a  greatly  increased  rate  of  evolution  of  hydrogen.f 
The  colour  of  the  hydrogen  flame  will  be  seen  to  change,  and 
to  assume  a  lilac  tint  (resembling  the  colour  given  to  a  flame  by 
potassium  compounds),  and  at  the  same  time  white  fumes  of  As4O6 
escape  from  the  tip  of  the  flame.  If  now  a  porcelain  dish  be  de- 
pressed upon  the  flame,  a  rich  brown-black  metallic-looking  stain 
will  be  deposited.  The  deposit  being  volatile,  and  the  flame  very 
hot,  the  stain  will  again  disappear  if  the  flame  be  allowed  to 
impinge  for  more  than  a  moment  or  two  on  the  same  spot. 

If  the  drawn-out  tube  be  heated  near  one  of  the  constrictions, 
the  arsenuretted  hydrogen  will  be  decomposed  as  it  passes  the  hot 
spot,  and  an  arsenic  mirror  will  be  deposited  in  the  tube. 


FIG.  13. 


*  A  small  test-tube  should  be  filled  by  upward  displacement,  and  tested  by 
a  flame  before  igniting  the  gas  at  the  exit-tube  of  the  apparatus.  As  an 
additional  precaution,  it  is  well  to  throw  a  duster  lightly  over  the  Woulffs 
bottle  before  applying  a  light,  so  that,  should  an  explosion  happen,  the  broken 
glass  will  be  prevented  from  flying  about. 

f  On  this  account,  it  is  necessary  that  the  generation  of  hydrogen  before 
adding  the  arsenic  solution  should  be  quite  slow  ;  and  also  that  the  quantity  of 
the  arsenic  solution  added  at  a  time  should  be  small. 


96  Qualitative  Analysis. 

It  will  be  noticed  that  the  deposition  takes  place  entirely  on 
that  part  of  the  tube  which  is  on  the  side  of  the  flame  farthest  from 
the  generating  vessel  *  (antimony  is  deposited  from  its  hydride  on 
both  sides  of  the  heated  spot). 

Since  antimony  also  forms  a  gaseous  compound  with  hydrogen 
which  gives  similar  stains,f  it  is  necessary  to  employ  further 
confirmatory  tests. 

1.  The  arsenic  stains  are  readily  dissolved  by  a  solution  of  a 
hypochlorite.      If,  therefore,  a   solution   of  bleaching   powder   be 
poured  over  such  stains  they  immediately  disappear — 

5Ca(OCl)2  +  6H9O  +  As4  =  5CaCl2  +  4H3AsO4 
Or- 

3Ca(OCl)2  +  2H2O  +  2AsH  =  sCaCl2  +  2H3AsO4 

Antimony  stains  do  not  dissolve  in  hypochlorite  solutions. 

2.  When  a  stream  of  H2S  is  passed  through  the  tube  containing 
the  deposit  of  either  arsenic  or  antimony,  slightly  warmed,  in  each 
case  the  sulphide  is  formed.     Yellow  arsenious  sulphide  volatilises 
along  from  the  warm  region  and  condenses  on  the  cold  distant  part 
of  the  tube  ;  antimonious  sulphide,  reddish  or  nearly  black,  remains 
unmoved,  being  non-volatile.     (If  present  together,  they  can  in  this 
way  be  separated.) 

If  a  stream  of  gaseous  HC1  be  now  passed  through  the  tube, 
antimonious  sulphide  is  converted  into  antimonious  chloride,  which 
passes  on  with  the  HC1,  and  may  be  led  into  water  and  again 
precipitated  as  the  red  sulphide  with  H2S.  The  yellow  arsenious 
sulphide  remains  in  the  tube,  being  unattacked  by  HC1. 

3.  Arsenuretted  hydrogen  can  also  be  distinguished  from  the 
antimony  compound,  by  the  difference  in  the  behaviour  of  the  two 
gases  towards  silver  nitrate.     When  passed  into  the  silver  solution, 
each  gas  produces  a  black  precipitate.     In  the  case  of  arsenic  this 
consists   of    metallic   silver,    while   with   antimony   it   consists  of 
antimonide  of  silver  ;  thus  — 

6AgN03  +  3H2O  +  AsH3  =  3Ag2  +  6HNO3  +  H3AsO3 
3AgN03  +  SbH3  =  SbAg3  +  3HNO3 

If  both  antimony  and  arsenic  were  originally  present,  these  two 

*  When  the  quantity  of  arsenic  present  is  very  minute,  it  will  not  be  visible 
in  the  flame,  neither  may  it  be  possible  to  obtain  a  stain  on  cold  porcelain. 
But  the  formation  of  the  mirror  in  the  heated  tube  is  a  method  by  which 
extremely  small  traces  of  arsenic  can  be  detected. 

f  The  stains  given  byantimoniuretted  hydrogen  have  a  rather  more  velvety 
or  sooty  appearance,  when  deposited  on  porcelain  from  the  flame,  than  those 
of  arsenic. 


Group  II. — Division  2.  97 

precipitates  will  be  produced  together.  On  filtering,  the  arsenic 
(now  as  arsenious  acid)  goes  into  the  nitrate  along  with  the  excess 
of  silver  nitrate  used.  Its  presence  may  be  detected  by  the  cautious 
addition  of  ammonia,  which  causes  the  precipitation  of  yellow  silver 
arsenite. 

The  antimony  (as  silver  antimonide)  remains  on  the  filter  ;  after 
being  washed,  it  is  boiled  with  a  solution  of  tartaric  acid.*  The 
liquid  thus  obtained  after  filtration  is  acidulated  with  HC1,  and 
antimonious  sulphide  precipitated  by  H2S. 

Fleitmanrfs  Test. 

When  an  arsenite,  or  a  solution  of  arsenious  oxide,  is  warmed 
in  a  test-tube  with  a  solution  of  sodium  hydroxide  and  metallic  zinc, 
arsenuretted  hydrogen  is  evolved,  which  can  be  detected  by  means 
of  a  piece  of  filter-paper  moistened  with  silver  nitrate  held  to  the 
mouth  of  the  tube.  A  black  stain  of  precipitated  silver  is  pro- 
duced. Antimoniuretted  hydrogen  is  not  produced  from  antimony 
compounds  under  similar  conditions. 

Antimony,  Sb. 

DRY  REACTIONS. — Antimony  compounds  may  be  reduced  to 
metallic  antimony  by  heating  them  with  Na2CO3  and  KCy  upon 
charcoal.  Globules  of  the  metal  are  thus  obtained,  which  burn  in 
the  blowpipe  flame,  producing  white  fumes  of  antimonious  oxide, 
Sb4O6,  the  combustion  being  continued  for  a  short  time  after 
removal  from  the  flame.  The  charcoal  at  the  same  time  receives  a 
white  incrustation.  The  bead  of  metal  will  be  found  to  be  very 
brittle,  and,  when  broken,  to  exhibit  a  highly  crystalline  appearance. 
Antimony  is  unacted  upon  by  dilute  HC1  or  H2SO4.  Nitric  acid 
oxidises  it  into  antimonic  acid,  or  antimonious  oxide,  depending 
upon  conditions  of  concentration. 

WET  REACTIONS. — Both  in  its  physical  properties  and  chemical 
relations  antimony  approaches  more  nearly  to  the  true  metals  than 
is  the  case  with  arsenic.  It  forms  two  series  of  compounds,  "anti- 
monious "  and  "  antimonic,"  which  may  be  regarded  as  being 
derived  respectively  from  the  two  oxides,  antimonious  oxide,  Sb4O6, 
and  antimony  pentoxide,  Sb2O5. 

*  Although  ordinary  metallic  antimony  is  not  soluble  in  tartaric  acid,  and  is 
only  slowly  attacked  by  strong  hydrochloric  acid  even  when  the  metal  is  in  the 
form  of  powder,  nevertheless,  when  combined  with  silver  as  it  is  in  the  precipitate 
of  silver  antimonide,  the  antimony  is  comparatively  easily  dissolved  by  HC1, 
and  dissolves  when  boiled  in  a  strong  solution  of  tartaric  acid.  In  the  latter 
case  antimony  tartrate,  (SbO)2C4H4p6,  goes  into  solution,  leaving  metallic 
silver.  With  HC1  antimonious  chloride  is  formed,  and  silver  chloride  remains. 

H 


98  Qualitative  Analysis. 

(a)  Antimonious  Compounds. — In  antimonious  oxide  we 
see  the  gradual  fading  away,  so  to  speak,  of  the  acidic  properties 
exhibited  by  the  corresponding  oxides  of  arsenic  and  phosphorus 
(elements  with  which  antimony  is  associated  in  the  natural  classi- 
fication), and  the  beginnings  of  basic  qualities.  Thus,  no  acids 
corresponding  to  this  oxide  are  known,  and  only  a  few  salts  (derived 
from  the  hypothetical  metantimonious  acid)  have  been  obtained. 
Of  these  the  best  known  is  sodium  metantimonite,  obtained  by 
dissolving  the  oxide  in  sodium  hydroxide  ;  thus — 

Sb4O6  +  4NaHO  =  4NaSbO2  +  2H2O 

On  the  other  hand,  this  oxide  unites  with  certain  acids  forming 
salts,  in  which  the  antimony — in  combination  with  oxygen  as  the 
monovalent  radical  antimony  I '(SbO) — takes  the  place  of  the  positive 
or  basilous  element.  Of  these  salts  the  tartrate,  (SbO)2(C4H4O6), 
and  the  double  potassium  tartrate  (tartar  £^#<r),(SbO)K(C4H4O6), 
are  the  most  familiar — 

Sb406  +  4HK(C4H406)  =  4(SbO)K(C4H400)  +  2H2O 

For  the  following  reactions  *  an  acid  (HC1)  solution  of  antimonious 
chloride  may  be  employed. 

KHO,  NaliO,  NH4HO,  as  well  as  alkaline  carbonates,  pre- 
cipitate antimonious  oxide,  Sb4O6  ;  thus — 

4SbCl3  +  I2KHO  =  Sb4OG  +  I2KC1  +  6H2O 
4SbCl3  +  6K2CO3  =  Sb4Oc  +  I2KC1  +  6CO2 

The  precipitate  redissolves  in  excess  of  either  potassium  or 
sodium  hydroxide,  forming  the  respective  metantimonites  (equation 
above). 

HO,  added  in  considerable  quantity  to  the  acid  solution  of 
SbCl3,  gives  a  white  precipitate  of  an  oxychloride,  SbOCl — 

SbCl3  +  H2O  =  2HC1  +  SbOCl 

This  compound  is  at  once  distinguished  from  the  similarly  pro- 
duced BiOCl,  by  the  fact  that  the  antimony  oxychloride  readily 
dissolves  in  tartaric  acid,  giving  antimonyl  tartrate  ;  thus — 

2SbOCl  +  H2(C4H406)  =  (SbO)2(C4H406)  +  2HC1 

On  boiling  the  precipitated  oxychloride  for  some  time  with  water, 
the  whole  of  its  chlorine  is  given  up,  and  antimonious  oxide  is 
formed — 

4SbOCl  +  2H2O  =  Sb4Oc  +  4HC1 

*  Except  that  with  AgNO3,  in  which  obviously  the  presence  of  chlorine 
would  interfere. 


Group  77.  —  Division  2.  99 

H2S  and  (NH4)2S  give  a  red  or  orange-red  precipitate  of  anti- 
monious  sulphide,  Sb2S3.  Soluble  in  excess  of  ammonium  sulphide. 
If  yellow  ammonium  sulphide  be  employed,  the  thio  salt  of  Sb2S5  is 
formed,  and  on  acidifying  the  liquid,  Sb2S5  is  precipitated  from  it. 
Antimonious  sulphide  is  soluble  also  in  caustic  alkalies,  from  which 
solution  the  trisulphide  is  again  thrown  on  acidifying — 

2Sb2S3  +  4KHO  =  3KSbS2  +  KSbO2  +  2H2O 
3KSbS2  +  KSb02  +  4HC1  =  2Sb2S3  +  4KC1  +  2H2O 

Antimonious  sulphide  is  not-dissolved  by  ammonia  or  by  ammonium 
carbonate  (contrast  arsenic). 

Antimonious  sulphide  is  decomposed  by  hot  hydrochloric 
acid — 

Sb2S3  +  6HC1  =  SbCl3  +  3H2S 

This  reaction  is  the  reverse  of  that  by  which  the  sulphide  is 
formed.  In  this  case,  however,  one  of  the  products,  namely 
the  H2S,  is  driven  from  the  sphere  of  action  as  fast  as  it  is  gene- 
rated. 

AgNO!}  added  to  an  alkaline  solution  of  antimonious  oxide, 
gives  a  black  precipitate,  consisting  of  silver  oxide  and  metallic 
silver.  The  action  takes  place  in  two  stages.  The  silver  nitrate 
interacts  with  the  caustic  alkali  present,  giving  silver  oxide,  and  this 
in  its  turn  gives  up  oxygen  to  the  antimon//*,  converting  it  into 
antimomz/tf  ;  thus — 

2AgNO3  +  2KHO  =  Ag2O  +  2KN03  +  H2O 
KSbO2  +  Ag2O  =  Ag2  +  KSbO3 

(b)  Antimonic  Compounds.  —  These  are  derived  from  the 
pentoxide,  Sb2O5.  This  oxide  is,  however,  so  feebly  acidic  that  it 
is  not  converted  into  antimonic  acid  by  the  action  of  water.  Only 
two  classes  of  antimonates  are  known  (compare  arsenates  and 
phosphates),  namely,  pyro-antimonates  (e.g.  K4Sb2O7)  and  met- 
antimonates  (e.g.  KSbO3). 

Potassium  pyro-antimonate  is  readily  soluble  in  water,  whiles 
the  sodium  salt  is  difficultly  soluble,  hence  the  potassium  compound 
is  used  as  a  reagent  for  sodium  (p.  22). 

KHO,  NaHO,  NH4HO,  as  well  as  alkaline  carbonates,  give 
with  an  acid  solution  of  antimony  pentachloride  a  white  precipitate 
consisting  of  metantimonic  acid — 

SbCl5  +  sKHO  =  HSbO3  +  jKCl  +  2H2O 


ioo  Qualitative  Analysis. 

H2O,  added  in  considerable  quantity,  produces  the  same  pre- 
cipitate * — 

SbCl5  +  3H20  =  HSb03  +  5HC1 

From  aqueous  solutions  of  antimonates,  the   same  precipitate  is 
produced  by  acids — 

KSbO3  +  HNO3  =  KNO3  +  HSbO3 

In  the  case  of  pyro-antimonates,  the  "  pyro  "  acid  is  first  pro- 
duced, which  then  passes  into  the  "  meta  "  compound  (footnote). 

H2S  and  (NH4)2S  produce  with  acidified  solutions  of  anti- 
monic  compounds  an  orange-red  precipitate  consisting  of  Sb2S5, 
Sb2S3,  and  S  in  varying  proportions. 

The  reducing  action  of  the  H2S'  converts  a  portion  of  the  anti- 
monic  compound  to  the  antimonious  state,  from  which  by  further 
action  the  antimonious  sulphide  is  thrown  down.  The  three  follow- 
ing reactions,  therefore,  proceed  simultaneously — 

2KSb03  +  5H2S  =  Sb2S5  +  5H20  +  K2O 

KSbO3  +  H2S  =  S  +  KSbO2  +  H2O 
2KSbO2  +  3H2S  =  Sb2S3  +  3H2O  +  K2O 

The  action  being  carried  on  in  acid  solution,  the  K^O  is  at  once 
neutralised. 

The  precipitate  behaves  towards  alkaline  sulphides  and  caustic 
alkalies  in  a  manner  similar  to  that  of  Sb2S3  ;  thus,  with  KHO 
potassium  thio-antimonate  (ortho)  and  antimonate  (meta)  are  pro- 
duced, from  which  an  acid  reprecipitates  the  Sb2S5 — 

4Sb2S5  +  iSKHO  =  5K3SbS4  +  3KSbO3  +  gH,O 
5K3SbS4  +  3KSbO3  +  iSHCl  =  4Sb2S5  +  iSKCl  +  9H2O 

AgNO3  gives  a  white  precipitate  with  aqueous  solutions  of  anti- 
monates, consisting  of  silver  antimonate.  The  precipitate  dissolves 
in  ammonia  (distinction  between  "  antimonic  "  and  "  antimonious  " 
compounds}. 

KI. — When  antimonic  compounds,  in  presence  of  HC1,  are 
boiled  with  a  solution  of  KI,  the  antimonic  compound  is  reduced 
to  the  "antimonious"  state,  and  iodine  is  liberated.  The  free 
iodine,  if  in  quantity,  colours  the  solution  brown  ;  if  in  small 
quantity,  it  may  be  recognised  by  the  formation  of  the  blue  colour 

*  This  action  takes  place  in  two  stages,  pyro-antimonic  acid  being  first 
formed,  which  loses  a  molecule  of  water  and  passes  into  the  more  stable 
metantimonic  acid ;  thus — 

2SbCl5  +  7H2O  =  loHCl  +  Sb2O5,2H2O(  =  H4Sb2O7) 
Sb2O5,2H2O  -  H2O  =  Sb2O5,H2O  =  2HSbO3 


Group  II.—  Division  2.  101 

with  starch  (distinction  between  "  antimonic"  and  (l antimonious  " 
compounds} — 

KSbO3  +  2HC1  +  2KI  =  KSbO2  +  2KC1  +  H2O  +  I2 
Or,  in  the  case  of  the  oxide  or  hydrated  oxide — 
Sb2O5,H2O  +  4HC1  +  4KI  =  Sb2O3,H2O  +  4KC1  +2H2O  +  2l2 

Precipitation  of  Metallic  Antimony.— Antimony  com- 
pounds (of  either  state  of  oxidation),  when  in  acid  (HC1)  solution, 
readily  deposit  metallic  antimony  by  galvanic  action.  The  test  is 
applied  in  the  following  way  :  one  or  two  drops  of  the  acid  solution 
are  placed  upon  a  piece  of  clean  platinum  foil,  and  a  small  frag- 
ment of  zinc  immersed  in  the  liquid.  Immediately  a  black  stain  is 
produced  upon  the  platinum  by  the  deposition  of  metallic  antimony. 
Even  very  dilute  solutions  give  the  stain,  hence  the  test  is  a  delicate 
one.  HC1  has  no  action  upon  the  deposit,  but  warm  HNO3  instantly 
attacks  it,  giving  antimonic  acid  (hence  the  absence  of  nitric  acid 
from  the  solution  before  applying  the  test  is  desirable). 

Antimoniuretted  hydrogen  is  evolved  when  antimony 
compounds  are  acted  upon  by  nascent  hydrogen.  The  compound 
undergoes  reactions  similar  to  those  of  the  corresponding  arsenic 
compound.  The  methods  for  distinguishing  between  them  are 
described  under  arsenic. 

Tin,  Sn. 

DRY  REACTIONS.— Compounds  of  tin  are  reduced  to  the  metallic 
state  by  being  heated  on  charcoal  with  Na2CO3  and  KCy  in  the 
reducing  flame.  A  portion  of  the  metal  is  oxidised  by  the  flame, 
and  produces  a  white  incrustation  of  SnO2  upon  the  charcoal ; 
this,  on  being  moistened  with  cobalt  nitrate  and  reheated,  assumes 
a  greenish  appearance. 

The  beads  of  reduced  metal  are  malleable  (therefore  easily  dis- 
tinguished from  Bi  or  Sb),  but  are  not  soft  enough  to  mark  paper 
in  the  manner  of  lead. 

Tin  dissolves  in  hot  strong  HC1,  forming  stannous  chloride, 
SnCl2,  hydrogen  being  evolved.  Hot  H2SO4  converts  it  into 
stannous  sulphate,  SnSO4,  and  gives  off  SO2. 

Cold  dilute  HNO3  dissolves  tin,  forming  stannous  nitrate, 
Sn(NO3)2,  and  ammonia  ;  ordinary  strong  acid  (sp.  gr.  1-24)  attacks 
it  violently,  converting  it  into  white  metastannic  acid,  H10Sn5O15. 

WET  REACTIONS. — Tin  forms  two  classes  of  compounds,  dis- 
tinguished as  "  stannous  "  and  "  stannic,"  derived  respectively  from 
stannous  oxide,  SnO,  and  stannic  oxide,  SnO2. 


IO2  Qualitative  Analysis. 

(a)  Stannous  Compounds. — Stannous  oxide  is  basic  in  its 
character,  giving  with  acids  the  Stannous  salts,  in  which  the  metal 
is  divalent.  Of  these  the  chloride,  sulphate,  and  nitrate  are  soluble 
in  water.  The  chloride  is  most  common. 

KHO,  NaHO,  NH4HO,  as  well  as  alkaline  carbonates,  give 
with  stannous  chloride  a  white  precipitate  of  hydrated  Stannous 
oxide  (basic  hydroxide)  ;  thus — 

2SnCl2  +  4KHO  =  2SnO,H2O  +  4KC1  +  H2O 

The  precipitate  is  soluble  in  excess  of  KHO  or  NaHO  (but  not  in 
the  other  precipitants),  forming  alkaline  stannites ;  thus — 

2SnO,H2O  +  4KHO  =  2K2SnO2  +  3H2O 

H2S  or  (NH4)2S  gives,  with  dilute  solutions  of  stannous  chloride, 
a  deep  brown  precipitate  of  stannous  sulphide.  Soluble  in  yellow 
ammonium  sulphide,  with  the  formation,  not  of  a  thto-stannite,  but 
thio-stannate.  (Colourless  ammonium  sulphide  is  almost  without 
action  upon  it.)  Stannic  sulphide,  SnS2,  is  thrown  down  upon  the 
addition  of  acids. 

Stannous  sulphide  dissolves  in  caustic  alkalies,  forming  a 
stannite  and  thio-stannite — 

2SnS  +  4KHO  =  K2SnO2  +  K2SnS2  +  2H2O 

From  this,  on  addition  of  acid,  stannous  sulphide  is  again  pre- 
cipitated— 

K2Sn02  +  K2SnS2  +  4HC1  =  4KC1  +  2H2O  +  2SnS 

Stannous  sulphide  is  insoluble  in  ammonium  carbonate.  Boil- 
ing HC1  converts  it  into  SnCl2 ;  while  aqua  regia  oxidises  it  to 
stannic  chloride,  SnCl4. 

Oxidation  of  Stannous  Compounds. — These  substances 
readily  pass,  by  oxidation,  into  stannic  compounds  ;  they  therefore 
act  the  part  of  powerful  reducing  agents  in  a  number  of  reactions, 
of  which  the  following  are  important  : — 

(1)  Mercuric  chloride,  in  the  presence  of  small  quantities  of 
stannous   chloride,   is   reduced   to   mercurous   chloride,  which   is 
thrown  down  as  a  white  precipitate.  With  excess  of  the  stannous  salt, 
and  on  warming,  the  precipitate  is  further  reduced,  and  becomes 
grey  through  the  separation  of  metallic  mercury — 

2HgCl2  +  SnCl2  =  Hg2Cl2  +  SnCl4 
Hg2Cl2  +  SnCl2  =  2Hg  +  SnCl4 

(2)  Ferric  salts  are  reduced  to  the  "  ferrous  "  state  :  thus, 


Group  II. — Division  2.  103 

ferric  sulphate,  in  the  presence  of  hydrochloric  acid,  gives  ferrous 
sulphate  and  stannic  chloride — 

Fe2(SO4)3  +SnCl2  +  2HC1  =  2FeSO4  +  H2SO4  +  SnCi; 

(3)  Gold   chloride,   in  the  presence  of  acid,  is  reduced  to 
metallic  gold — 

2AuCl3  +  3SnCl2  =  3SnCl4  +  2Au 

In  neutral  solutions,  a  reddish  or  purple  precipitate  (or  colora- 
tion, if  very  dilute)  is  produced,  known  as  purple  of  Casstus.  Its 
formation  is  promoted  by  the  presence  of  a  little  stannic  chloride 
in  the  stannous  compound.  Its  composition  is  believed  to  be  ex- 
pressed by  the  formula  Au2,3SnO2.  Acids  convert  it  into  metallic 
gold  and  a  stannic  salt. 

(4)  Cupric    salts,    chromates,  and  permanganates  are 
reduced  respectively  to  cuprous,  chromic,  and  manganous  salts  ; 
thus— 

2KMnO4  +  i6HCl  +  sSnCla=  2KC1  +  2MnCl2  +  8H2O  +  5SnCl4 

(5)  Bismuth  salts  are  reduced  in  alkaline  solutions,  with  the 
precipitation  of  black  bismuthous  oxide,  Bi2O2  or  BiO,  and  the 
oxidation  of  the  stannite  to  stannate.    Thus,  if  bismuth  chloride  or 
nitrate  (salts  derived  from  Bi2O3)  are  added  to  an  alkaline  solution 
of  stannous  oxide  (i.e.  to  a  solution  of  potassium  stannite),  the 
following  reaction  takes  place  : — 

K2SnO2  +  6KHO  +  2Bi(NO3)3  =  K2SnOs  +  6KNO3  +  3H2O  + 

2BiO 

The  reaction  will  be  simpler  if  we  have  regard  only  to  the  oxide 
of  bismuth,  from  which  the  nitrate  is  derived — 

K2SnO2  +  Bi2O3  =  2BiO  +  K2SnO3 

(6)  An  aqueous  solution  of  stannous  chloride  gradually  absorbs 
atmospheric  oxygen,  being  converted  partly  into  stannic  chloride, 
and  a  white  insoluble  basic  chloride  ;  thus — 

3SnCl2  +  H2O  +  O  =  SnCl4  +  SnCl2,SnO,H2O 

(b)  Stannic  Compounds. — Stannic  oxide,  SnO2,  may  be 
regarded  as  being  both  a  "  basic  "  and  an  acidic  oxide,  for,  although 
itself  insoluble  in  either  acids  or  alkalies,  we  may  consider  both 
the  stannic  salts  and  the  stannates  as  being  derived  from  this 
oxide.  The  most  important  stannic  salt  is  the  chloride,  SnCl4.  A 
solution  of  this  compound  in  hydrochloric  acid  may  be  used  for 
the  following  reactions  : — 


IO4  Qualitative  Analysis. 

KHO,  NaHO,  NH4HO,  as  well  as  alkaline  carbonates,  give 
a  white  precipitate  of  hydrated  stannic  oxide,  or  stannic  acid, 
SnO2,H2O,  or  H2SnO3*  ;  thus— 

SnCl4  +  4NaHO  =  H2SnO3  +  4NaCl  +  H2O 

Soluble  in  HNO3  and  in  HC1.     Soluble  in  KHO  and  NaHO,  with 
formation  of  the  respective  stannates,  K2SnO3,  and  Na2SnO3 ;  thus— 

H2SnO3  +  2NaHO  =  Na2SnO3  +  2H2O 

H2S  or  (NH4)2S  precipitates  yellow  stannic  sulphide,  SnS2 
(with  H2S  the  precipitate  appears  nearly  white  at  first,  and  is  only 
complete  in  dilute  solutions).  The  precipitate  is  soluble  in  caustic 
alkalies,  in  ammonium  sulphide,  and  sulphides  of  the  alkalies, 
forming  thio-stannates  ;  thus — 

3SnS2  +  6KHO  =  2K2SnS3  +  K2SnO3  +  sH2O 
SnS2  +  (NH4)2S  =  (NH4)2SnS3 

From  these  solutions  the  yellow  stannic  sulphide  is  reprecipitated 
on  the  addition  of  HC1— 

(NH4)2SnS3  +  2HC1  =  2NH4C1  +  SnS2  +  H2S 
Stannic  sulphide  is  insoluble  in  ammonium  carbonate,  but  dis- 
solves in  hot  strong  HC1.  Stannic  oxysalts,  such  as  the  nitrate  or 
sulphate,  are  unstable  in  aqueous  solution,  and  are  decomposed 
into  stannic  or  metastannic  acid.  Therefore,  when  by  double 
decomposition  such  oxysalts  might  be  expected  to  form,  the  result 
<is  the  precipitation  of  one  or  both  of  these  stannic  acids.  Thus, 
when  sulphuric  acid  is  added  to  stannic  chloride  and  the  solution 
diluted  with  water,  the  action  may  be  represented  by  the  two 
equations — 

(1)  SnCl4  +  2H2S04  =  4HC1  +  Sn(SO4)2 

(2)  [Sn(S04)2+  3H20  =  2H2S04  +  H2SnO3]  x  5  to  re- 

present metastannic  acid. 

A  similar  precipitation  takes  place  when  neutral  salts  of  the 
alkalies  are  employed,  such  as  sodium  sulphate  or  ammonium  i 
nitrate;  thus  — 

*  Meta-stannic  acid,  the  white  compound  obtained  by  the  action  of  nitric 
acid  upon  tin,  is  expressed  by  the  same  formula  multiplied  by  five,  5(H2SnO3), 
or  H10Sn5Oi5.  It  forms  salts  by  the  replacement  of  two  hydrogen  atoms 
only,  as  is  the  case  with  stannic  acid  ;  their  composition  may  therefore  be  ex- 
pressed by  the  formula  (taking  potassium  metastannate  as  an  example), 
K2bnO3,4SnO2(4H2O.  Boiling  (or  fusing)  with  caustic  alkalies  converts  meta- 
sJannates  into  stannates ;  thus — 

K2SnO3,4SnO2)4H2O  -f  8KHO  =  sK2SnO3  +  8H2O 


Group  II. — Division  2.  105 

(1)  SnCl4  +  4(NH4)N03  =  4NH4C1  +  Sn(NO3)4 

(2)  [Sn(N03)4  +  3H20  =  4HNO3  +  H2SnO3]  x  5 

Precipitation  of  Metallic  Tin. — When  zinc  is  immersed  in 
an  acid  solution  of  stannous  or  stannic  chloride,  the  tin  is  displaced 
by  the  zinc,  and  precipitated  as  a  grey-black  deposit  upon  the 
surface  of  the  zinc  ;  or,  if  the  whole  of  the  zinc  becomes  dissolved, 
the  tin  is  left  as  a  scaly  powder.  It  may  be  collected,  and,  after 
being  washed,  dissolved  in  hot  hydrochloric  acid.  Tin  produces 
no  stain  upon  platinum,  as  is  the  case  with  antimony  ;  and,  again, 
differs  from  this  metal  in  dissolving  in  HC1. 

Gold  and  Platinum. 

Although  these  two  metals  do  not  belong  to  the  class  of  rare 
elements,  nevertheless  it  is  very  rarely  that  the  student  is  called 
upon  to  analyse  mixtures  containing  either  of  them.  In  actual 
practice  these  metals  are  met  with  only  in  the  analysis  of  alloys, 
where  their  presence  is  probably  more  than  suspected  at  the  outset. 
Neither  gold  nor  platinum  are  soluble  in  either  sulphuric,  nitric, 
or  hydrochloric  acid  ;  but  they  readily  dissolve  in  a  mixture  of 
nitric  and  hydrochloric  acids,  yielding  the  chlorides  AuCl3  and 
PtCl4  respectively.  Comparatively  few  simple  salts  containing 
gold  or  platinum  are  known,  and  their  compounds  generally  are 
characterised  by  the  extreme  readiness  with  which  they  are  reduced 
to  the  metallic  state. 

H2S  or  (NH4)2S  gives,  with  solutions  of  AuCl3  and  PtCl4,  pre- 
cipitates of  the  sulphides.  In  the  case  of  platinum  the  precipita- 
tion is  slow.  Auro-auric  sulphide,  Au2S,Au2S3,  or  AuS  (black), 
and  platinic  sulphide,  PtS2  (also  black) — 

8AuCl3  +  9H2S  +  4H2O  =  24HCl  +  H2SO4  +  2(Au2S,Au2S3),  orSAuS 
PtCl4  +  2H2S  =  4HC1  +  PtS2 

If  the  gold  solution  be  boiling,  the  reduction  goes  further,  and 
metallic  gold  is  thrown  down — 

8AuCl3  +  3H2S  +  I2H2O  =  24HC1  +  3H2SO4  +  8Au 

The  sulphides  of  both  metals  are  insoluble  in  HC1  and  HNO3, 
but  dissolve  in  aqua  regia. 

AuS  dissolves  in  ammonium  sulphide  and  in  sulphides  of  the 
alkalies,  forming  thio-aurates.  PtS2,  unmixed  with  other  sulphides, 
does  not  dissolve  in  ammonium  sulphide  ;  but  in  the  presence  of 
other  sulphides  of  the  group,  it  is  partially  dissolved.*  From  the 

*  Owing  to  the  fact  that  PtS2  is  only  partially  dissolved  by  ammonium 
sulphide,  if  platinum  happened  to  be  present  in  a  mixture  which  was  under- 
going systematic  analysis,  a  portion  of  the  sulphide  would  pass  into  Group  II. 
Division  2,  along  with  As,  Sb,  Sn,  Au ;  and  a  part  of  it  would  remain  in 
Division  i,  along  with  Hg,  Bi,  Cu,  Cd,  being  ultimately  found  with  the  HgS, 
insoluble  in  HNO3.  For  this  reason  it  is  more  advantageous  to  remove 
platinum  (and  also  gold)  before  the  separation  of  Group  II.  is  commenced, 
aS  explained  on  p.  107. 


106  Qualitative  Analysis. 

solution  in  both  cases,  the  sulphides  are  reprecipitated  on  the 
addition  of  HC1. 

KHO. — Neither  the  hydroxides  nor  carbonates  of  the  alkalies 
give  any  precipitate  with  moderately  dilute  solutions  of  AuCl3  or 

Jtr  t  \—  1^  • 

From  a  concentrated  solution  of  AuCl3,  NH4HO  gives  an 
orange-red  precipitate  of  fulminating  gold,  (NH3)2,Au2O3  ;  while 
KHO  produces  a  brown  precipitate  of  hydrated  auric  oxide, 
Au2O3,3H2O,  or  Au(HO)3.  This  precipitate  is  soluble  in  excess  of 
potash,  yielding  potassium  aurate,  K2O,Au2O3,  or  KAuO2. 

Both  AuCl3  and  PtCl4  form  double  salts  with  alkaline  chlorides 
(chloro-aurates  and  chloro-platinates).  Ammonium  chloro-aurate, 
NH4Cl,AuCl3,  or  NH4AuCl4,  is  soluble  in  water  ;  while  ammonium 
chloro-platinate,  2NH4Cl,PtCl4,  or  (NH4)2PtCl6,  is  moderately 
insoluble,  and  is  produced  by  precipitation  when  ammonium 
chloride  is  added  to  a  moderately  strong  solution  of  PtCl4. 

Gold  and  platinum  are  readily  reduced  from  their  compounds 
(more  especially  gold)  and  precipitated  in  the  metallic  State, 
and  their  most  characteristic  reactions  are  based  upon  this  fact* 
Thus— 

1.  Ferrous  sulphate  gives  a  brown  precipitate  of  metallic  gold ; 
in  weak  solutions  a  bluish  coloration — 

AuCl3  +  3FeSO4  =  Fe2(SO4)3  +  FeCl3  +  Au 

With  platinum  the  action  only  takes  place  on  prolonged  boiling— 

3PtCl4  +  i2FeSO4  =  4Fe2(SO4)3  -f  4FeCl3  +  3?t 

2.  Oxalic  acid,  on  being  gently  warmed  with  AuCl3,  causes  the 
deposition   of   the   metal   either  as   a  scaly  precipitate   or  as   a 
coherent  gold  film  upon  the  glass,  according  to  the  conditions  of 
the  experiment.     Platinum  is  not  reduced  by  oxalic  acid — 

2AuCl3  +  3C2H204  =  6C02  +  6HC1  +  2Au 

3.  Potassium  nitrite  reduces  AuCl3,  being  itself  converted  into 
nitrate — 

2AuCl3  +  3H2O  +  3KNO2  =  3KNO3  +  6HC1  +  2Au 

With  platinum  no  precipitate  forms  at  first,  but  on  standing,  yellow 
crystals  are  deposited  of  a  double  nitrite  of  potassium  and  platinum 
(one  of  the  few  oxysalts  of  platinum  known)  containing  platinum  in 
\heplatmous  condition,  2KNO2,Pt(NO2)2,  or  K2Pt(NO2)4*— 

PtCl4  +  5KNO2  +  H2O  =  K2Pt(NO2)4  +  2KC1  +  2HC1  +  KN03 

4.  Stannous  chloride  gives  with  AuCl3  a  precipitate  or  colora- 
tion   (depending   upon    concentration)  varying    in    colour    from 
reddish-brown  to  purple.     The  compound  is  known  as  purple  of 

*  The  platino-nitrites  are  remarkable  in  that  the  platinum  they  contain 
does  not  answer  to  the  ordinary  tests  for  that  metal,  just  as  the  iron  in  ferro- 
cyanides  is  not  detected  by  the  ordinary  reactions  for  iron. 


Separation  of  the  Metals  of  Group  II. — Division  2.   107 

Cassius,  and  its  composition  is  not  known  with  certainty.  The 
presence  of  a  small  quantity  of  stannic  chloride  (such  as  is  always 
present  in  a  solution  of  stannous  chloride  except  when  quite 
freshly  made)  facilitates  the  production  of  the  purple. 

With  PtCl4  a  brown  colour  is  produced,  by  the  reduction  of  the 
platinum  chloride  to  platinous  chloride,  PtCl2. 

In  analysis,  when  gold  and  platinum  are  present,  it  is  prefer- 
able to  remove  them  before  the  precipitation  of  Group  II.  by 
sulphuretted  hydrogen.  The  gold  is  precipitated  in  the  metallic 
state  by  oxalic  acid,  and  the  solution  is  evaporated  down  with 
ammonium  chloride,  which  causes  the  precipitation  of  the  platinum 
as  ammonium  platinum  chloride,  2NH4Cl,PtCl4. 


SEPARATION  OF  THE  METALS  OF  GROUP  II.— DIVISION  2. 

The  separation  of  Group  II.  from  Groups  III.,  IV.,  and  V. 
depends  upon  the  precipitation  of  their  sulphides  from  acid 
solutions. 

The  separation  of  Division  i  from  Division  2  is  based  upon 
the  solubility  of  the  sulphides  of  the  latter  in  ammonium  sulphide 
or  in  caustic  alkalies. 

The  separation  of  the  metals  As,  Sb,  and  Sn  from  each  other  is 
based  upon — 

(1)  The  insolubility  of  arsenic  sulphide  in  hydrochloric  acid  (or 
its  solubility  in  ammonium  carbonate),  whereby  arsenic  is  separated 
from  Sb  and  Sn. 

(2)  The  solubility  of  metallic  tin,  and  insolubility  of  antimony, 
in  hydrochloric  acid. 

The  solution,  if  neutral  or  alkaline,  is  acidified  with  HC1,*  and 
sulphuretted  hydrogen  passed  through  (as  described  on  p.  88)  until 
the  precipitation  of  the  metals  of  Group  II.  is  complete.  The  pre- 
cipitate is  thoroughly  washed  (in  order  to  make  the  separation  from 
Groups  III.,  IV.,  and  V.  complete),  and  is  then  transferred  to  a 
small  beaker,  and  gently  warmed  with  yellow  ammonium  sulphide 
for  a  few  minutes.f  The  liquid  is  then  filtered.  The  residue  con- 
sists of  the  undissolved  sulphides  of  the  metals  of  Group  II., 
Division  i. 


*  If  the  solution  under  examination  is  alkaline,  it  may  contain  thio  salts  of 
As,  Sb,  or  Sn  ;  the  addition  of  HC1  will  result  in  the  precipitation  of  the  sul- 
phides of  these  metals.  If  it  is  neutral,  basic  salts  of  antimony  might  be  pre- 
cipitated at  first,  but  redissolve  on  warming  with  a  slight  excess  of  the  acid. 

t  Ammonium  sulphide  dissolves  CuS  to  a  slight  extent  (see  Reactions,  p.  85), 
hence,  if  this  element  is  present,  a  small  quantity  of  it  will  find  its  way  into  the 
solution  along  with  As,  Sb,  and  Sn. 


io8  Qualitative  Analysis. 


The  solution  contains  the  thio  salts  of  As,  Sb,  and  Sn.  It  should  be  some- 
what diluted,  and  hydrochloric  acid  added  drop  by  drop  until  the  sul- 
phides are  completely  reprecipitated ;  then  filtered  and  washed.  The 
precipitate  is  then  transferred  to  a  boiling-tube  with  a  small  quantity 
of  HC1,  and  boiled  for  a  few  moments  until  H2S  is  no  longer  given 
off.  It  is  then  diluted  and  filtered. 


The  residue  consists  of 
arsenic  sulphide  and 
sulphur. 

Confirm  by  dissolving  in 
HC1  with  a  crystal  of 
KC1O3,  and  applying 
special  reactions  for 
arsenic,  such  as  Fleit- 


man's  or  Reinsch's  test. 


The  solution.  Pour  a  few  drops  upon  a  piece 
of  platinum  foil,  and  add  a  fragment  of  zinc. 
A  black  stain  indicates  Sb.  If  antimony  is 
present,  place  a  strip  of  zinc  along  with  the 
platinum  foil  in  the  remainder  of  the  solution, 
until  all  the  antimony  and  tin  are  thrown 
down.  Collect  the  deposit  and  boil  it  with 
strong  HC1,  and  filter.  Test  the  filtrate  for 
Sn  by  means  of  HgCU. 


APPENDIX   TO    CHAPTER   IX. 

THE  RARE  METALS  OF  GROUP  II. 

Four  of  these  metals  belong  to  Division  i  of  this  group,  their 
sulphides  being  insoluble  in  ammonium  sulphide  ;  these  are — 

Ruthenium,  Rhodium,  Palladium,  and  Osmium. 

The  remaining  members  form  sulphides  which  are  soluble  in 
ammonium  sulphide,  and  they  therefore  belong  to  the  antimony, 
arsenic,  and  tin  subdivision,  namely — 

Iridium,  Tellurium,  Selenium,  Molybdenum. 

The  composition  of  the  precipitates  which  are  thrown  down  by 
the  group-reagent,  is  the  following  : — 
Ru2S3j 

p  ,  A  3>  Insoluble  in  ammonium  sulphide. 

OsS    j 

Ir2S3  j 

ce  2  [Soluble  in  ammonium  sulphide, 
oe 

MoS3J 

*  In  order  to  ascertain  the  condition  of  oxidation  in  which  the  arsenic 
originally  existed  in  the  substance  under  analysis,  special  tests  must  be  applied 
to  the  solution  before  it  has  been  exposed  either  to  reducing  or  oxidising 
influences,  as  in  the  case  of  iron. 


The  Rare  Metals  of  Group  II.  109 

The  metals  of  the  first  division,  together  with  iridium  in  the 
second  section,  belong  to  the  natural  family  of  elements  known  as 
the  platinum  metals,  because  they  all  occur  associated  together  in 
platinum  ore.  Of  these,  ruthenium  and  rhodium  are  the  most 
rare.* 

Palladium. — This  metal,  in  the  elemental  state,  is  readily  dis- 
tinguished from  all  the  others  of  the  platinum  group  by  its  ready 
solubility  in  warm  nitric  acid.  The  other  platinum  metals  are 
unacted  upon  by  any  ordinary  acid.  Aqua  regia  is  without  action 
upon  rhodium  and  iridium;  it  acts  with  slowness  upon  ruthenium, 
and  readily  dissolves  platinum,  forming  the  chloride,  while  it  con- 
verts osmium  into  the  tetroxide. 

When  palladium  is  dissolved  in  nitric  acid,  the  compound  formed 
is  palladious  nitrate,  Pd(NO3)2.  If  the  solution  be  diluted  with 
water,  especially  if  the  amount  of  free  acid  present  is  only  small,  a 
brown-coloured  precipitate  is  produced,  consisting  of  a  basic 
nitrate.  Palladious  sulphide,  PdS,  produced  by  the  group  reagent, 
is  black. 

Mercuric  cyanide,  HgCy2,  gives  a  yellowish  precipitate  of  palla- 
dious cyanide,  PdCy2  ;  slightly  soluble  in  hydrochloric  acid,  readily 
soluble  in  ammonia  and  in  potassium  cyanide.  The  precipitate  is 
distinguished  from  other  metallic  cyanides  by  the  reaction  common 
to  all  palladium  salts,  namely,  that  when  heated  they  decompose, 
leaving  spongy  metallic  palladium. 

Potassium  iodide,  KI,  gives  a  characteristic  black  precipitate  of 
palladious  iodide,  PdI2.  Palladious  salts  are  readily  reduced  to  the 
metallic  state  either  by  heat  or  by  the  action  of  reducing  agents. 

Osmium. — The  compound  of  this  rare  element  which  is  most 
commonly  met  with  is  the  so-called  osmic  acid,  which  is  employed 
in  the  preparation  of  microscopic  sections  of  animal  tissues.  This 
compound  is  the  tetroxide,  OsO4,  or  osmic  anhydride. 

It  is  characterised  by  its  extremely  low  melting-point  (about  40°) 
and  boiling-point  (100°),  and  by  the  peculiar  and  irritating  vapour 
which  it  gives.  The  vapour  exerts  a  most  injurious  effect  upon  the 
eyes,  and  is  extremely  poisonous. 

This  vapour  is  given  off  when  any  osmium  compound  is  heated 
with  nitric  acid,  and  serves  as  a  characteristic  test  for  the  element. 

The  tetroxide  is  soluble  in  water,  giving  a  neutral  solution  which 
has  powerful  oxidising  properties  ;  it  bleaches  indigo,  liberates 
iodine  from  potassium  iodide,  and  oxidises  ferrous  sulphate  and 
alcohol,  the  osmium  compound  being  reduced  to  the  state  of 
hydrated  dioxide,  OsO2,2H2O  (or  Os(HO)4),  which  is  thrown  down 
as  a  black  precipitate. 

Sulphurous  acid,  or  a  sulphite  added  to  the  solution,  produces  a 
series  of  colour-changes  from  yellow  to  green  and  lastly  blue,  the 
colour  of  the  osmious  sulphite,  OsSO3,  which  then  separates  out. 

Iridium. — This  metal  differs  from  platinum  in  not  being  dis- 
solved by  aqua  regia.  When  heated  with  a  fused  mixture  of  sodium 

*  See  footnote  on  p.  70. 


no  Qualitative  Analysis. 

nitrate  and  hydroxide  in  a  silver  vessel,  the  metal  is  oxidised  to  the 
trioxide,  Ir2O3  ;  and  on  treating  the  residue  with  aqua  regia,  a 
dark-coloured  solution  is  obtained  of  the  double  chloride.  2NaCl, 
IrCl4. 

This  solution  may  be  used  for  the  following  reactions  : — 

When  sulphuretted  hydrogen  is  passed  into  the  solution,  the 
brown  colour  disappears  owing  to  the  reduction  of  the  IrCl4  to 
IrCl3  (or  Ir2Cl6)  and  simultaneous  precipitation  of  sulphur.  The 
further  passage  of  the  gas  throws  down  the  trisulphide,  Ir2S3,  as  a 
dark  brownish  precipitate. 

The  double  chlorides,  2NH4Cl,IrCl4  and  2KCl,IrCl4,  are  pre- 
cipitated by  the  addition  of  ammonium  chloride  and  potassium 
chloride  respectively.  They  are  both  dark  brownish-red  precipitates, 
insoluble  in  strong  solutions  of  the  precipitants. 

By  the  action  of  reducing  agents  (e.g.  ferrous  or  stannous  salts, 
nitrites,  etc.),  these  double  chlorides  are  reduced,  giving  similar 
compounds  containing  the  lower  chloride  of  indium,  and  having 
the  composition  expressed  by  the  formulae  3NH4Cl,IrCl3  and 
3KCl,IrCl3  respectively.  The  solution  is  at  the  same  time  decolor- 
ised, and  the  double  chloride  gradually  deposits.  When  caustic  alkali 
is  added  to  a  solution  of  iridic  chloride  (or  the  double  sodium  salt), 
and  the  mixture  heated,  the  solution  assumes  a  deep  blue  colour 
owing  to  the  precipitation  of  iridic  hydroxide,  Ir(HO)4,  which  when 
separated  appears  as  an  indigo-blue  powder.  This  reaction  serves 
to  distinguish  iridium  from  platinum. 

Tellurium  and  Selenium. — These  two  elements  belong  to 
the  same  natural  family  as  sulphur,  which  in  some  respects  they 
closely  resemble.  They  both  lie  on  the  borderland  between  the 
non-metals  and  the  true  metals.  While  selenium  forms  no  stable 
compounds  in  which  it  forms  the  positive  constituent,  tellurium 
exhibits  feeble  basic  properties  ;  the  oxide,  TeO2,  being  both  an 
acid-forming  and  a  salt-forming  oxide. 

Both  elements  form  hydrogen  compounds  corresponding  to 
sulphuretted  hydrogen,  and  closely  resembling  it  in  properties. 
Thus,  when  compounds  of  either  element  are  heated  on  charcoal 
with  sodium  carbonate,  sodium  telluride,  Na2Te,  or  selenide, 
Na2Se,  is  formed.  These  are  decomposed  by  acids,  with  forma- 
tion of  the  respective  hydrogen  compounds,  whose  odour  is  even 
more  offensive  than  that  of  sulphuretted  hydrogen.  Or  if  the 
sodium  compounds  are  moistened  with  water  upon  a  silver  coin, 
a  black  stain  is  produced  of  silver  telluride  or  selenide. 

Sulphuretted  hydrogen  gives,  with  tellurous  compounds,  a  brown 
precipitate  of  tellurous  sulphide,  TeS2,  but  with  selenious  compounds 
the  precipitate  consists  of  selenium  and  sulphur,  selenious  sulphide 
being  too  unstable  to  exist.  Both  precipitates,  however,  are  soluble 
in  ammonium  sulphide.  With  selenzV  compounds  (selenates)  sul- 
phuretted hydrogen  gives  no  precipitate  until  the  selenate  is 
reduced  to  selenite. 

Barium  chloride,  BaCl2,  gives,  with  solutions  of  selenates,  a 
white  precipitate  of  barium  selenate,  BaSeO4.  This  precipitate  is 


The  Rare  Metals  of  Group  II.  1 1 1 

distinguished  from  barium  sulphate  in  that,  when  boiled  with  hydro- 
chloric acid,  it  is  converted  into  barium  selenzte,  which  is  soluble. 

Sulphurous  acid  reduces  both  tellurous  and  selenious  com- 
pounds, with  precipitation  of  the  element ;  tellurium  being  thrown 
down  as  a  black  powder,  while  selenium  is  precipitated  in  the  form 
of  the  brick-red  amorphous  variety. 

Both  elements  burn  in  the  air  or  in  oxygen,  with  a  blue 
flame,  giving  rise  to  the  dioxide ;  in  the  case  of  selenium,  the 
combustion  is  accompanied  by  a  smell  of  putrid  horseradish. 

Molybdenum. — The  compound  of  this  element  most  commonly 
met  with  is  ammonium  molybdate,  (NH4)2MoO4. 

From  a  strong  aqueous  solution  of  this  salt,  hydrochloric  or 
nitric  acid  gives  a  white  precipitate  of  molybdic  acid,  H2MoO4, 
soluble  in  excess  of  acid. 

When  sulphuretted  hydrogen  is  passed  into  an  acidulated  solu- 
tion, the  solution  first  assumes  a  blue  colour,  which  turns  green  as 
the  dark-brown  sulphide  is  precipitated.  The  precipitate  produced 
is  the  trisulphide,  MoS3,  analogous  to  the  trioxide.  It  is  soluble 
in  ammonium  sulphide  (and  alkaline  sulphides  generally),  forming 
thio-molybdates.  These,  like  the  corresponding  antimony  and 
arsenic  compounds,  are  decomposed  by  hydrochloric  acid,  with  the 
reprecipitation  of  the  tri-sulphide. 

The  reaction  by  which  molybdic  acid  is  most  readily  identified, 
is  the  formation  of  the  yellow  precipitate  of  ammonium  phospho- 
molybdate,  by  adding  sodium  phosphate  to  a  nitric  acid  solution 
of  molybdic  acid. 


CHAPTER  X. 
REACTIONS  OF  THE  METALS  OF  GROUP  I. 

Silver,  Lead,  Mercury  (as  Mercurous  Mercury). 

THE  classification  of  these  three  elements  into  a  group  is  owing  to 
common  properties  possessed  by  their  chlorides.  In  most  other 
respects  they  are  greatly  dissimilar,  and  in  the  "  natural "  classifi- 
cation of  the  elements  they  take  their  places  in  three  different 
families. 

Silver,  Ag. 

DRY  REACTIONS. — Compounds  of  silver,  when  heated  on  char- 
coal with  sodium  carbonate  in  the  reducing  flame,  yield  metallic 
silver  ;  which,  being  non-oxidisable,  is  not  accompanied  by  any 
oxide  incrustation  upon  the  charcoal.  The  metal,  however,  is 
slightly  volatile  in  the  blowpipe  flame,  and  sometimes  a  faint  red- 
brown  incrustation  is  thus  obtained. 

The  reduced  metal  may  be  removed  to  a  watch-glass,  dissolved 
in  nitric  acid,  and  precipitated  as  chloride. 

WET  REACTIONS. — The  salts  of  silver  are  derived  from  the 
monoxide  Ag2O.  Of  the  common  salts,  the  nitrate  is  readily 
soluble,  the  acetate  and  sulphate  sparingly  soluble,  in  water.  (The 
chlorate,  nitrite,  and  fluoride  are  also  soluble.  The  solubility  of 
the  fluoride  is  noteworthy  in  view  of  the  insolubility  of  the  other 
halogen  salts.) 

A  characteristic  property  exhibited  by  a  number  of  the  silver 
salts  is  their  readiness  to  form  soluble  compounds  with  ammonia, 
which  are  either  double  salts  of  silver  and  ammonium,  such  as 
NH4NO3,AgNO3,  or  belong  to  the  class  of  salts  known  as  metall- 
ammonium  compounds,  or  metallo-amines.  Thus,  silver  nitrate 
absorbs  ammonia  gas  and  yields  the  compound  AgNO3,2NH3. 
The  same  substance  is  obtained  on  adding  ammonia  solution  to 
silver  nitrate  (a  neutral  salt)  until  the  precipitate  first  formed  is 
dissolved. 


Group  L  113 

HC1,  and  soluble  chlorides,  give  a  white  amorphous  precipitate 
of  silver  chloride,  AgCl,  which,  on  being  warmed  or  stirred,  becomes 
granulated  in  appearance,  and  very  quickly  settles.  On  exposure 
to  light,  the  white  compound  assumes  a  slate  colour  or  drab  tint, 
which  gradually  deepens  to  a  violet,  and  finally  appears  brown  or 
black. 

Silver  chloride  is  quite  insoluble  in  water,  but  soluble  to  a 
slight  extent  in  strong  HC1 ;  dilution  causes  the  complete  precipi- 
tation. It  readily  dissolves  in  ammonia,  forming  the  compound 
2AgCl,3NH3.  Nitric  acid  decomposes  this  compound,  causing 
the  reprecipitation  of  AgCl,  which  is  practically  insoluble  in  that 
acid — 

2AgCl,3NH3  +  3HN03  =  3NH4NO3  +  2AgCl 

Silver  chloride  is  soluble  also  in  KCy,  being  first  converted  into 
silver  cyanide,  which  dissolves  in  excess  of  KCy,  forming  the  double 
cyanide  KCy,AgCy.  It  also  dissolves  in  sodium  thiosulphate,  with 
the  formation  of  a  double  thiosulphate  ;  thus  — 

AgCl  +  Na2S2O3  =  NaCl  +  NaAgS2O3 

When  boiled  with  potassium  hydroxide,  silver  chloride  is  con- 
verted into  silver  oxide  (black) — 

2AgCl  +  2KHO  =  2KC1  +  H20  +  Ag2O 

Silver  chloride  melts  without  decomposition  at  451°,  and  re- 
solidifies to  a  horny  mass  (horn  silver). 

[Reactions  with  bromides,  iodides,  and  cyanides  are  described 
under  the  respective  acids.] 

KHO,  NaHO,  or  NH4HO  gives  a  greyish-black  precipitate  of 
silver  oxide,  Ag2O.  Insoluble  in  excess  of  the  caustic  alkalies,  but 
readily  soluble  in  ammonia.  If  the  silver  solution  is  acid,  ammonia 
gives  no  precipitate,  but  forms  a  soluble  double  salt  (see  above). 

K2CO3  or  Na2CO3  gives  a  white  precipitate  of  silver  carbonate, 
Ag2CO3.  Insoluble  in  excess  of  the  precipitant ;  soluble  in  am- 
monium carbonate,  ammonia,  and  nitric  acid. 

H2S  or  (NH4)2S  produces  a  black  precipitate  of  silver  sulphide, 
Ag2S.  Insoluble  in  dilute  acids,  except  boiling  dilute  nitric  acid, 
which  converts  it  into  nitrate.  The  H2S,  which  by  double  de- 
composition is  set  free,  is  acted  upon  by  the  nitric  acid,  with  the 
precipitation  of  sulphur  and  evolution  of  nitric  oxide  (see  Lead 
reactions,  p.  80). 

Silver  sulphide  is  insoluble  in  ammonia,  ammonium  sulphide,  or 
potassium  sulphide. 

I 


U4  Qualitative  Analysis. 

Reduction  of  Silver  Salts  to  the  Metallic  State. — 

Silver  compounds  (more  especially  the  ammoniacal  solution  of 
silver  oxide)  are  readily  reduced  with  precipitation  of  metallic  silver 
(which  often  deposits  as  a  coherent  mirror)  by  certain  organic 
substances,  as  sugar,  tartrates,  aldehydes,  etc.  Many  inorganic 
salts  also,  which  act  as  reducing  agents,  precipitate  metallic  silver 
from  solutions  of  its  salts,  e.g.  ferrous  sulphate — 

3AgN03  +  3FeS04  =  Fe(NO3)3  +  Fe2(SO4)3  +  3Ag 

Many  metals  are  capable  of  reducing  silver  compounds.  Men- 
tion may  be  made  of  the  three  metals,  zinc,  iron,  and  mercury. 

If  a  strip  of  zinc  be  immersed  in  a  solution  of  silver  nitrate, 
crystals  of  metallic  silver  are  seen  to  grow  out  from  the  surface  of 
the  zinc  in  the  manner  of  the  familiar  "  lead  tree."    The  reducing 
action  of  zinc  is  often  employed  in  the  laboratory  to  convert  pre- 
cipitated silver  chloride  into  metallic  silver.     The  same  action  of 
iron  is  used  in  one  of  the  metallurgical  processes  for  the  extraction 
of  silver  ;  the  reaction  in  both  cases  is  similar — 
2AgCl  +  Zn  =  ZnCl2  +  2Ag 
2AgCl  +  Fe  =  FeCl2  +  2Ag 

Mercury  in  contact  with  silver  chloride  similarly  precipitates 
metallic  silver,  with  formation  of  mercurous  chloride — 
2AgCl  +  2Hg  =  Hg2Cl2  +  2Ag 

None  of  these  reductions  are  made  use  of  in  qualitative  analysis ; 
the  latter  reaction,  however,  with  mercury  has  an  important  bearing 
upon  the  separation  of  silver  from  mercurous  mercury  in  the  usual 
course  of  analysis.  The  separation  of  these  metals  is  based  upon 
the  solubility  of  silver  chloride  in  ammonia,  mercurous  chloride 
being  at  the  same  time  converted  into  the  black  mercurous 
ammonium  chloride — 

(i)  Hg2Q2  +  2NH4HO  =  NH2(Hg2)Cl  +  NH4C1  +  2H2O 

This  mercurous  compound,  however,  in  the  presence  of  silver 
chloride  and  excess  of  ammonia,  passes  into  the  corresponding  and 
more  stable  mercuric  salt,  NH2HgCl,  while  the  mercury  it  thus 
loses,  reduces  a  portion  of  the  silver  chloride  ;  thus — 

(2)  NH2(Hg2)Cl  +  2AgCl  =  NH2HgCl  +  HgCl2  +  2Ag* 

(3)  HgCl2  +  2NH4HO  =  NH2HgCl  +  NH4C1  +  2H2O 

*  It  is  possible,  but  not  probable,  that  the  following  reaction  goes  on 
simultaneously  : — 

2NH2(Hg2)Cl  +  2AgCl  =  2NH2HgCl  +  Hg2Cl2  +  2Ag 

The  HgjClai  being  reconverted  by  the  excess  of  ammonia  into  mercurous 
ammonium  chloride,  would  then  be  in  a  position  to  react  upon  a  fresh  pro- 
portion of  silver  chloride,  so  that  the  action  would  be  continued  until  the  whole 
of  the  silver  became  reduced  ;  this,  however,  does  not  appear  to  be  the  case. 


Separation  of  the  Metals  of  Group  I.          115 

The  entire   change,  therefore,  may  be  summed  up  in  the  single 
equation  following  :  — 


In  the  practical  separation,  therefore,  a  small  quantity  of  silver, 
in  presence  of  a  large  proportion  of  mercury,  might  escape  detection 
by  being  entirely  precipitated  and  left  upon  the  filter  along  with  the 
mercury  compound.  It  will  be  evident  that  if  the  action  of  the 
ammonia  be  allowed  to  continue  sufficiently  long,  it  would  be 
possible  to  precipitate  the  whole  of  the  silver  if  the  quantity  of 
mercury  present  only  slightly  exceeded  the  equivalent  proportion. 

Lead,  Mercury. 

The  reactions  of  these  metals  have  already  been  considered  in 
connection  with  the  metals  of  Group  II.,  Division  i,  pp.  76,  79. 

SEPARATION  OF  THE  METALS  OF  GROUP  I. 

The  separation  of  these  metals  from  the  other  groups,  depends 
upon  the  insolubility  of  their  chlorides  in  cold  water,  and  con- 
sequently their  precipitation  by  hydrochloric  acid.  The  separation 
of  the  metals  from  each  other  is  based  upon  — 

(1)  The  solubility  of  lead  chloride  in  hot  water. 

(2)  The  solubility  of  silver  chloride  in  ammonia. 

To  the  solution  add  moderately  dilute  hydrochloric  acid  drop 
by  drop,  until  a  slight  excess  beyond  what  is  required  for  complete 
precipitation  has  been  added.  Gently  warm  the  mixture,*  and 
after  again  cooling  it,  filter.  The  filtrate  contains  Groups  II.,  III., 
IV.,  and  V. 


The  precipitate,  consisting  of  PbCl?,  AgCl,  and  Hg2Cl2,  is  thoroughly 
washed  in  cold  water,  and  then  boiled  with  water  (or  washed  while  in 
the  filter  with  boiling  water)  and  filtered. 


The  filtrate  contains 
PbCl2,  which  deposits 
in  white  needle-shaped 
crystals  on  cooling. 

Confirm  by  special  test,  e.g. 
the  formation  of  PbCrO4. 


The  residue  is  treated,  while  still  upon  the 
filter,  with  a  small  quantity  of  ammonia, 
which  dissolves  the  AgCl,  and  converts  the 
white  Hg2Cl2  into  black  NH2(Hg2)Cl.t 

The  solution  yields  a  white  precipitate  of 
AgCl  upon  being  acidified  with  HNO3. 

The  black  residue  may  be  dissolved  in  a  little 
aqua  regia,  and  (after  being  nearly  neutralised) 
the  mercury  precipitated  upon  metallic  copper. 


*  See  footnote  on  p.  107. 

t  For  conditions  under  which  this  separation  is  incomplete,  see  p.  114. 


n6  Qualitative  Analysis. 


SYSTEMATIC  SEPARATION  OF  THE  GROUPS.* 

(1)  Separation  of  Group  I.— By  precipitation  with  hydro- 
chloric acid  (p.  115). 

The  precipitate  is  thoroughly  washed  with  cold  water,  and 
examined  by  the  method  on  p.  115. 

(2)  Separation  of  Group  II.— The  filtrate  from  Group  I.  is 
diluted  with  water  (any  precipitation  of  basic  bismuth  compounds 
is  to  be  disregarded,  as  the  subsequent  action  of  SH2  will  convert 
them  into  the  sulphide),  and  a  moderately  slow  stream  of  sulphu- 
retted hydrogen  passed  through.     The  changes  which  occur  during 
this   precipitation   should    be   carefully   watched    and   noted  (see 
Lead,  Mercury,  Arsenic  reactions).     Complete  precipitation  must 
be  ensured  (p.  88.     See  also  Arsenic,  p.  93).     The  precipitate  is 
then  treated  as  described  on  p.  107. 

(3)  Separation  of  Groups  IIlA.  and  IIlB.— The  filtrate 
from  Group  1 1.  is  boiled  until  the  sulphuretted  hydrogen  is  entirely 
expelled.     Two  or  three  drops  of  strong  nitric  acid  are  added,  and 
the  liquid  again  boiled  for  a  few  minutes,  in  order  to  oxidise  com- 
pounds (e.g.  iron  or  chromium)  which  might  have  become  reduced 
by  the  action  of  sulphuretted  hydrogen.   The  liquid  is  then  carefully 
evaporated  to  dryness  f  in  a  porcelain  dish,  and  if  it  shows  signs 

*  It  may  not  be  out  of  place  at  this  point  to  impress  once  more  upon 
the  student  the  supreme  importance  of  cleanliness,  neatness,  and  method  in 
analysis.  Slipshod  and  careless  work  inevitably  ends  in  disappointment  and 
failure.  All  the  utensils  (test-tubes,  beakers,  etc.)  must  be  scrupulously  clean, 
being  always  rinsed  once  or  twice  with  distilled  water  after  having  been  washed. 
Too  much  care  cannot  be  paid  to  securing  complete  precipitation  in  every 
separation,  or  to  the  thorough  washing  of  precipitates.  The  time  spent  in 
securing  these  results  is  never  lost,  whereas  the  slovenly  neglect  of  these  points 
may,  and  frequently  does,  involve  a  repetition  of  the  analysis,  and  a  corre- 
sponding sacrifice  of  time.  A  little  experience  will  enable  the  student  to  keep 
two  or  three  operations  going  on  at  the  same  time.  Thus,  while  one  precipitate 
is  being  washed,  another  precipitation  can  be  made,  and  the  filtration  of  this 
second  separation  can  be  carried  on  simultaneously  with  the  other.  While 
these  two  are  filtering,  other  tests  or  special  reactions  can  be  made.  To  do 
this  successfully,  however,  it  must  be  done  methodically,  and  (especially  in 
cases  where  the  analysis  has  to  be  interrupted  for  a  time)  the  various  precipi- 
tates and  solutions  should  be  labelled.  Before  making  a  group  separation,  it 
is  well  to  make  a  preliminary  test  in  a  small  portion  of  the  solution,  and  if  the 
group-reagent  gives  no  precipitate,  the  operation  may  then  be  omitted  with 
the  bulk  of  the  solution. 

f  "  Carefully  evaporate  to  dryness"  does  not  mean  that  the  student  may 
place  a  lamp  under  the  dish  and  go  away  and  leave  it.  The  operation  must 
be  watched,  and  as  the  liquid  becomes  more  and  more  concentrated  the  lamp 
flame  must  be  lowered.  As  the  residue  begins  to  dry  round  the  edges,  he 
should  notice  if  it  shows  signs  of  charring,  for  if  there  is  no  organic  compound 
present  it  is  better  to  avoid  strongly  heating  the  dried  residue,  as  the  operation 


Systematic  Separation  of  the  Groups.  1 1 7 

of  charring  (owing  to  the  presence  of  citrates,  tartrates,  etc.),  the 
residue  is  cautiously  heated  until  such  organic  matter  is  completely 
decomposed.  The  residue  is  moistened  with  a  few  drops  of  strong 
hydrochloric  acid,  water  is  then  added,  and  the  mixture  boiled. 
The  remaining  residue,  consisting  of  silica  and  carbon,  is  filtered 
off.* 

A  small  portion  of  the  solution  is  next  tested  for  phosphoric 
acid  by  means  of  ammonium  molybdate  (p.  63).  To  the  main 
portion  of  the  solution  ammonium  chloride  is  added  in  consider- 
able quantity,  and  the  liquid  heated  to  boiling.  Ammonia  is 
carefully  added  until  precipitation  is  complete. 

If  phosphoric  add  is  absent,  this  precipitate  is  examined  for 
the  metals  of  Group  III  A.  by  the  method  on  p.  49. 

If  phosphoric  add  is  present,  the  precipitate  is  treated  as  explained 
on  p.  68.  The  filtrate  is  then  saturated  with  sulphuretted  hydro- 
gen,! gently  warmed,  and  filtered.  The  precipitate  is  examined  for 
the  metals  of  Group  1 1  IB.  according  to  the  plan  on  p.  62. 

(4)  Separation  of  Group  IV. — The  filtrate  from  Group  1 1  IB. 
is  boiled  briskly,  with  the  addition  of  a  little  hydrochloric  acid, 
to  decompose  ammonium  sulphide  and  expel  all  the  sulphuretted 
hydrogen.    Any  precipitated  sulphur  is  removed  by  filtration.    The 
solution  is  then  rendered  alkaline  by  addition  of  ammonia,  and 
ammonium  carbonate  added  until  precipitation  is  complete.     The 
solution  may  be  warmed,  but  must  not  be  boiled  (see  p.  34).    The' 
precipitate  is  examined  for  the  metals  of  Group  IV.,  as  indicated 
on  p.  35. 

(5)  Detection  of  Metals  of  Group  V.— The  filtrate  from 
Group  IV.  is  tested  for  magnesium,  potassium,  and  sodium,}  as 
described  on  p.  25. 

is  likely  to  render  certain  oxides  (e.g.  Fe2O3,  Cr2O3,  A12O3)  very  difficult  of 
solution  in  hydrochloric  acid. 

*  If  by  a  preliminary  test  upon  a  small  portion  of  the  solution  (by  evapo- 
rating it  upon  a  platinum  capsule)  it  is  found  that  there  is  no  silicious  or 
carbonaceous  residue,  the  main  bulk  of  the  liquid  need  not  be  evaporated 
down. 

f  Or  ammonium  sulphide  may  be  used  for  the  precipitation  (see,  in  this 
case,  the  action  of  ammonium  sulphide  on  nickel  sulphide). 

J  Ammonium  obviously  cannot  be  tested  for  in  this  solution,  since  ammo- 
nium salts  have  been  frequently  introduced  during  the  course  of  analysis.  The 
test  for  this  "metal"  must  therefore  be  made  in  a  portion  of  the  original 
solution. 


1 1 8  Qualitative  A  nalysis. 


APPENDIX    TO     CHAPTER    X. 

THE  RARE  METALS  OF  GROUP  I. 

The  rare  metals  of  this  group  are  thallium  and  tungsten. 

The  compounds  that  are  precipitated  by  the  group-reagent 
being  thallous  chloride,  T1C1,  and  tungstic  acid,  H2WO4,  in  a 
hydrated  condition. 

Thallium. — Two  classes  of  compounds  of  thallium  are  known, 
namely,  thallous  salts  (derived  from  thallous  oxide,  T12O)  and 
thallic  salts  (derived  from  thallic  oxide,  T12O3).  In  the  thallous 
compounds  (which  are  the  more  stable  of  the  two  classes),  the 
element  shows  a  resemblance  to  the  alkali  metals  on  the  one  hand, 
and  to  lead  on  the  other.  For  example,  thallous  hydroxide,  T1HO, 
is  soluble  in  water,  giving  a  strongly  alkaline  solution  ;  hence 
KHO,  NaHO,  and  NH4HO  give  no  precipitate  with  thallous 
solutions.  Thallous  carbonate,  T12CO3  is  moderately  soluble  in 
water,  and  therefore  is  only  precipitated  by  carbonates  of  the  alka- 
lies from  concentrated  solutions  (compare  Thallic  salts).  Again, 
platinic  chloride  gives  with  thallous  solutions  an  orange-yellow 
precipitate  of  thallous  platinic  chloride,  2TlCl,PtCl4.  On  the 
other  hand,  its  resemblance  to  lead  is  specially  seen  in  the  chloride, 
iodide,  and  chromate. 

Hydrochloric  acid  (or  soluble  chlorides)  gives  a  white  curdy 
precipitate  of  T1C1.  Like  lead  chloride  it  is  slightly  soluble  in 
cold  water,  and  dissolves  more  readily  in  hot  water  (100  parts 
of  water  at  16°  dissolve  0-265  parts,  while  at  100°  1*427  parts 
are  dissolved). 

It  is  at  once  distinguished  from  lead  chloride  by  the  fact 
that  it  is  readily  soluble  in  strong  sulphuric  acid,  forming  soluble 
thallous  sulphate.  It  is  distinguished  from  silver  chloride  in  not 
being  soluble  in  ammonia,  and  in  not  changing  colour  on  exposure 
to  light.  Potassium  iodide  gives  a  bright  golden-yellow  precipitate 
of  thallous  iodide,  Til  ;  while  potassium  chromate  throws  down 
thallous  chromate,  Tl2CrO4,  also  (like  the  lead  compound)  yellow 
in  colour. 

Thallous  sulphide  (black)  is  readily  soluble  in  mineral  acids  ; 
hence  sulphuretted  hydrogen  only  gives  complete  precipitation  in 
acetic  acid  solutions  (or  in  alkaline  solutions).  With  ammonium 
sulphide  precipitation  is  complete,  but  the  precipitate  rapidly  under- 
goes atmospheric  oxidation  into  the  sulphate,  which  passes  into 
solution.  Thallic  salts  are  readily  distinguished  from  thallous 
compounds.  Thus  the  chloride,  T1C13,  is  soluble  in  water  ;  hence 
hydrochloric  acid  gives  no  precipitate.  Potassium  chromate,  simi- 
larly, gives  no  precipitate.  With  potassium  iodide,  a  mixture  of 
thallous  iodide  and  iodine  is  precipitated,  while  with  sulphuretted 
hydrogen  the  compound  is  reduced  to  the  thallous  state,  with 
precipitation  of  sulphur. 


The  Rare  Metals  of  Group  I.  119 

The  caustic  alkalies,  and  also  the  alkaline  carbonates,  give  a 
brown  precipitate  with  lhallic  solutions,  consisting  of  an  oxy- 
hydroxide,  TIO(HO). 

Thallium,  like  lead,  is  easily  reduced  from  solutions  of  its  salts 
by  metallic  zinc,  the  thallium  being  deposited  upon  the  zinc  in  a 
spongy  metallic  state. 

When  introduced  into  a  Bunsen  flame  upon  a  platinum  wire, 
thallium  salts  impart  a  characteristic  brilliant  green  colour,  which, 
when  viewed  through  a  spectroscope,  is  seen  to  consist  of  one 
bright  green  line. 

Tungsten. — The  compound  of  this  element  most  commonly 
met  with  in  commerce  is  sodium  tungstate.  The  formula  repre- 
senting the  composition  of  the  normal  $2\\.  is  Na2WO42H2O.*  The 
tungstates  of  the  alkalies  are  alone  soluble.f  Hydrochloric  acid 
gives  a  white  precipitate  of  hydrated  tungstic  acid.  When  air-dried, 
this  precipitate  has  the  composition  WO(HO)4,  or  H2WO4,H2O. 
When  dried  over  sulphuric  acid,  it  loses  H2O  and  is  converted 
into  the  normal  acid.  The  precipitate  is  insoluble  in  excess  of 
acid. 

The  most  characteristic  reaction  for  compounds  of  tungsten  is 
the  blue  colour  given  when  either  metallic  zinc  or  tin  or  stannous 
chloride  is  added  to  a  solution  of  a  tungstate,  and  the  solution 
strongly  acidified  with  hydrochloric  acid.  On  the  addition  of 
stannous  chloride  a  pale  yellow  precipitate  is  produced,  which,  on 
the  addition  of  hydrochloric  acid  and  gently  warming,  turns  deep 
blue.* 

Sulphuretted  hydrogen  gives  no  precipitate  in  acid  solutions, 
but  causes  the  solution  to  assume  a  blue  colour.^ 

Tungsten  trisulphide,  WS3,  is  a  thio-anhydride,  forming  soluble 
thio-tungstates  with  alkaline  sulphides  ;  hence,  ammonium  sulphide 
gives  no  precipitate  with  solutions  of  alkaline  tungstates.  But  on 
acidulating  the  mixture  with  hydrochloric  acid,  the  sulphide  is 
thrown  down  as  a  dark  brown  precipitate. 

*  There  is  quite  a  number  of  sodium  tungstates,  which  may  all  be  regarded 
as  compounds  of  the  normal  salt  with  varying  quantities  of  the  trioxide.  Of 
these  the  so-called  metatungstate  and  paratungstate  are  articles  of  commerce — 

Sodium  metatungstate,  Na2W4Oi3,ioH2O  ;  or  Na2WO4,3WO3,ioH2O 
Sodium  paratungstate,  Nai0Wi2O4i,2i(or  28)H2O  ; 

or  5Na2WO4,7WO3,2i(or  28)H2O 

The  former  of  these  two  salts  is  the  common  compound  sometimes  used  for 
rendering  fabrics  uninflammable. 

f  The  oxides  of  tungsten  are  all  acidic  in  their  characters,  and  no  salts  of 
this  element  are  known  in  which  it  functions  as  the  positive  or  basic  constituent. 

%  The  blue  colour  is  believed  to  be  due  to  the  formation  of  a  lower  hydrated 
oxide. 


CHAPTER  XI. 

THE  NON-METALS  AND  THEIR  ACIDS. 

ONE  of  the  chief  chemical  distinctions  between  metals  and  non- 
metals  is  that,  while  the  oxides  of  the  former  exhibit,  generally 
speaking,  basic  properties,  the  latter  elements  give  oxides  which 
are  acidic  in  their  character.  Two  exceptions  to  this  generalisation 
respecting  the  non-metals  are  seen  in  the  case  of  hydrogen  (whose 
oxides  are  not  acidic),  and  the  element  oxygen  itself.  Although 
the  number  of  non-metals  is  comparatively  very  small,  the  number 
of  acids  derived  from  them  is  very  considerable,  owing  to  the  fact 
that  many  of  these  elements  give  rise  to  several  acids. 

All  the  elements  belonging  to  this  section  form  compounds  with 
hydrogen  ;  but  only  in  the  case  of  fluorine,  chlorine,  bromine, 
iodine,  and  sulphur  are  the  hydrogen  compounds  included  among 
the  acids  in  analytical  classification. 

The  classification  of  the  acids  (or  acid  radicals)  is  based,  as  in 
the  case  of  the  metals,  upon  the  solubility  or  insolubility  of  certain 
of  their  salts  produced  by  interaction  with  certain  specified  reagents. 
They  are  in  this  way  divided  into  certain  arbitrary  groups,  but  the 
mode  of  treatment  differs  entirely  from  that  adopted  in  the  case  of 
the  metals.  The  reagents  referred  to  are  not  employed  as  "  group- 
reagents  "  to  separate  one  group  of  acid  radicals  from  another ; 
but  each  acid  is  separately  detected  by  a  special  test.  The  use  of 
the  general  reagent  is  in  order  to  ascertain  by  a  single  operation 
the  absence,  or  otherwise,  of  an  entire  group,  whereby  the  necessity 
for  applying  a  number  of  separate  tests  may  be  obviated. 

The  study  of  the  special  reactions  for  the  non-metals  and  their 
acids  may  therefore  be  made  irrespective  of  their  analytical 
classification,  and  it  will  be  more  advantageous  to  postpone  the 
consideration  of  the  classification  until  a  knowledge  of  the  special 
reactions  has  been  gained,  and  the  student  is  prepared  to  undertake 
the  systematic  detection  of  the  acids. 


The  Non-metals  and  their  Acids.  12 1 

The  non-metals  and  their  acids  which  will  be  included  in  this 
section  are  the  following  : — 

Chlorine :  hydrochloric  acid,  hypochlorous  acid,  chloric  acid, 
perchloric  acid. 

Bromine  :  hydrobromic  acid,  bromic  acid. 

Iodine :  hydriodic  acid,  lodic  acid. 

Fluorine  :  hydrofluoric  acid,  hydrofluo-silicic  acid. 

Sulphur :  sulphuretted  hydrogen,  sulphuric  acid,  sulphurous 
acid,  thio-sulphuric  acid. 

Nitrogen  :  nitric  acid,  nitrous  acid. 

Phosphorus  :  phosphoric  acid,  phosphorous  acid,  hypophos- 
phorous  acid. 

Carbon :  carbonic  acid,  formic  acid,  oxalic  acid,  acetic  acid, 
tartaric  acid,  citric  acid,  cyanogen,  hydrocyanic  acid,  ferrocyanic 
acid,  ferricyanic  acid,  cyanic  acid,  thiocyanic  acid. 

Silicon :  silicic  acid. 

Boron :  boric  acid. 

In  some  instances  the  properties  of  the  elements  themselves,  in 
the  elemental  state,  are  made  use  of  in  analysis  ;  such,  for  instance, 
as  in  the  cases  of  chlorine,  bromine,  and  iodine.  When  this  is  the 
case,  those  properties  of  the  elements  by  means  of  which  they  are 
most  readily  identified  will  be  studied.  With  others,  such  as 
fluorine,  silicon,  boron,  the  properties  of  the  isolated  elements 
have  no  analytical  bearing,  and  therefore  all  description  of  such 
elements  will  be  omitted.* 

Certain  acids,  such  as  arsenious,  arsenic,  chromic,  have  already 
been  discussed  under  their  respective  metals. 

The  Halogens. 

These  elements  not  only  form  oxy-acids,  but  they  also  combine 
with  hydrogen  and  yield  acids.  The  salts  of  these  hydrogen  acids, 
e.g.  chlorides,  bromides,  and  iodides  are  known  as  the  haloid  salts, 
or  sometimes  as  the  halides.  • 

Chlorine. 

The  properties  by  which  this  element  is  identified  are  the 
following : — 

It  is  a  pale  greenish-yellow  gas,  having  a  characteristic  suffocat- 
ing smell,  irritating  and  rapidly  attacking  the  mucous  membrane 
of  the  nose  and  throat.  It  dissolves  in  water,  imparting  its  own 

11  It  is  presupposed  that  the  student  of  analysis  has  already  become  familiar 
with  the  common  properties  of  the  non-metallic  elements  ;  and  if  he  has  not, 
they  must  be  sought  in  text-books  of  general  chemistry. 


122  Qualitative  Analysis. 

colour  to  the  solution.  It  combines  readily  with  many  metals  ;  it 
is  a  powerful  oxidising  agent ;  it  liberates  bromine  and  iodine  from 
bromides  and  iodides  respectively  ;  it  possesses  powerful  bleaching 
properties. 

The  Liberation  of  Chlorine  from  its  Compounds. — The 
method  usually  resorted  to  in  analysis  for  the  liberation  of  chlorine, 
depends  upon  the  action  of  peroxides  (manganese  dioxide  being 
employed)  upon  the  hydrogen  acid,  or  upon  a  chloride  in  the 
presence  of  sulphuric  acid — 

2NaCl  +  2H2SO4  +  MnO2  =  Na2SO4  +  MnSO4  +  2H2O  +  C12 

For  this  purpose  the  mixture  of  the  chloride  with  sulphuric  acid 
and  manganese  dioxide  is  gently  warmed  in  a  small  flask  or  test- 
tube,  fitted  with  a  delivery  tube,  and  the  evolved  gas  is  allowed  to 
pass  into  water  contained  in  a  second  test-tube.  As  the  solution 
of  the  chlorine  by  water  is  not  complete,  the  characteristic  smell  of 
that  which  escapes  solution  may  be  noted,  and  its  bleaching  action 
can  be  seen  by  introducing  a  strip  of  litmus  paper  into  the  mouth 
of  the  tube. 

The  presence  of  the  free  chlorine  in  the  water  may  be  detected 
by  the  following  more  delicate  tests  : — 

(a)  On  adding  a  few  drops  of  the  chlorine  water  to  a  solution 
of  potassium  iodide,  to  which  a  little  dilute  starch  paste  has  been 
added,  a  deep  blue  coloration  results,  owing  to  the  liberated  iodine 
(set  free  by  the  chlorine)  uniting  with  the  starch. 

(b]  A  crystal  of  ferrous  ammonium  sulphate*  [FeSO4,(NH4)2- 
SO4,6H2O]  is  dissolved  in  water,  and  a  few  drops  of  ammonium 
thiocyanate  added.    To  this  colourless  mixture  a  drop  or  two  of  the 
chlorine  water  is  added.     This  at  once  oxidises  the  ferrous  salt  to 
the  ferric   state,  which,  in  the   presence  of  the  thiocyanate,  im- 
mediately   gives   rise    to   the  wine-red  coloration    due  to    ferric 
thiocyanate. 

Hydrochloric  Acid  and  Chlorides. 

Hydrochloric  acid  is  a  colourless  gas  having  a  sharp  choking 
smell.  It  fumes  in  contact  with  moist  air,  is  strongly  acid,  but  has 
no  bleaching  properties.  It  is  extremely  soluble  in  water,  the 
solution  constituting  the  ordinary  reagent. 

Hydrochloric    acid  is   liberated    in  the    gaseous    state  when 

*  This  double  salt  is  used  instead  of  ferrous  sulphate,  as  it  is  less  easily 
oxidised  by  the  air,  and  therefore  is  more  easily  obtained  free  from  ferric 
compounds. 


Hydrochloric  Acid  and  Chlorides.  123 

chlorides  (except  those  of  tin,  lead,  mercury,  and  silver)  are  gently 
heated  with  strong  sulphuric  acid  — 

KC1  +  H2S04  =  HKSO,  +  HC1 

The  presence  of  free  hydrochloric  acid  in  a  solution  containing 
a  soluble  chloride  may  be  detected  by  gently  warming  the  liquid 
with  manganese  dioxide  (without  the  addition  of  sulphuric  acid)  ; 
chlorine  is  evolved,  which  may  be  detected  by  the  methods  already 
described  — 

4HC1  +  MnO2  =  MnCl2  +  2H2O  +  C12 

Chlorides  are  all  soluble  in  water,  except  those  of  the  metals  of 
Group  I.  (PbCl2  being  soluble  in  hot  water),  and  certain  others 
which  are  decomposed  by  water.  The  chief  of  these  latter  salts 
are  the  chlorides  of  (phosphorus,  arsenic),  antimony,  bismuth,  and 
tin.  In  contact  with  water  these  give  either  an  oxide  or  an 
oxychloride  of  the  metal,  with  formation  of  free  hydrochloric  acid 
(see  Reactions  of  these  metals). 

Silver  nitrate,  AgNO3,  gives,  in  solution  of  chlorides  or 
hydrochloric  acid,  a  white  precipitate  of  silver  chloride.  Insoluble 
in  nitric  acid.  Readily  soluble  in  ammonia,  even  dilute  (for 
further  properties,  see  Silver  reactions,  p.  113). 

AgCl  is  distinguished  from  either  AgBr  or  Agl  by  the  fact 
that  chlorine  water  is  without  action  upon  it  (see  Bromides  and 
iodides}.  It  may  also  be  distinguished  in  the  following  way  :  — 

Reduction  by  Zinc.  —  If  the  washed  precipitate  of  AgCl  be 
mixed  with  a  little  very  dilute  sulphuric  acid,  and  a  strip  of  zinc 
placed  in  the  mixture,  the  silver  chloride  turns  grey,  owing  to  its 
reduction  to  metallic  silver,  while  zinc  chloride  passes  into  solution. 
This,  on  treatment  with  manganese  dioxide  and  sulphuric  acid,  will 
yield  chlorine. 

Fusion  with  sodium  carbonate  converts  AgCl  into  metallic 
silver  and  sodium  chloride.  On  treatment  with  water,  chlorine  can 
be  liberated  from  the  solution,  as  in  the  foregoing. 

Formation  of  Chromyl  Chloride.—  When  a  chloride  is  mixed 
with  potassium  dichromate  (the  two  salts  being  powdered  together), 
and  the  mixture  gently  warmed  with  strong  sulphuric  acid,  a  red- 
brown  vapour  is  disengaged  (resembling  bromine  in  colour,  but  very 
different  in  smell)  consisting  of  chromyl  chloride,  CrO2Cl2  — 

4KC1  +  K2Cr2O7  +  3H2SO4  =  3K2SO4  +3H2O  +  2CrO2Cl2  * 


*  We  may  regard  this  action  as  taking  place  between  chromium  trioxide 
(formed  by  the  action  of  the  sulphuric  acid  upon  the  dichromate)  and  hydro- 
chloric acid  (from  the  interaction  of  the  chloride  and  sulphuric  acid)  ;  thus  — 
CrO3  +  2HC1  =  H2O  +  CrO2Cl2 


124  Qualitative   Analysis. 

The  gas  is  decomposed  by  water  or  alkaline  hydroxides,  form- 
ing in  the  latter  case  an  alkaline  chromate  and  chloride ;  thus — 
CrO2Cl2  +  4NH4HO  =  (NH4)2CrO4  +  2NH4C1  +  2H2O 

If  the  reaction  be  made  in  a  test-tube  or  small  flask  fitted  with  a 
delivery  tube,  and  the  vapour  of  the  chromyl  chloride  (which  is  a 
deep  red  fuming  liquid  at  ordinary  temperatures)  be  passed  into  a 
second  test-tube  containing  ammonia,  the  above  decomposition 
takes  place.  The  presence  of  ammonium  chromate  is  seen  by  the 
yellow  colour  which  the  liquid  assumes,  and  the  presence  of  the 
chromate  is  proof  of  the  presence  of  a  chloride  in  the  first  test-tube 
or  flask.  [The  presence  of  the  chromate  may  be  further  confirmed 
by  applying  a  special  test,  such  as  the  reaction  with  hydrogen 
peroxide.] 

By  means  of  this  test  it  is  possible  to  detect  a  chloride  in  the 
presence  ^/either  a  bromide  or  iodide,*  as  neither  bromine  nor  iodine 
form  similar  chromyl  compounds.  When  a  bromide  or  iodide  is 
treated  in  this  way,  bromine  or  iodine  vapour  escapes  alone. 

Bromine. 

Bromine  is  a  dark,  brown-red,  volatile  liquid,  which  passes  into 
a  brown-red  vapour  at  ordinary  temperatures.  It  has  a  powerful, 
irritating  smell,  attacks  the  mucous  membrane  of  the  nose  and 
throat,  and  causes  the  eyes  to  smart.  It  bleaches,  but  less  easily 
than  chlorine.  It  produces  a  yellow  colour  with  starch.  It  dis- 
solves in  water,  giving  a  reddish  solution  (bromine  water],  and  is 
soluble  in  ether  and  in  carbon  disulphide,  giving  reddish-brown 
solutions. 

Liberation  from  its  Compounds. — Bromine  is  liberated 
from  bromides  by  the  action  of  manganese  dioxide  and  sulphuric 
acid  (compare -chlorine).  The  solution  of  the  bromide  and  the 
manganese  dioxide  are  placed  in  a  small  beaker,  and  a  little  strong 

It  is  important  to  bear  in  mind  that  with  an  excess  of  hydrochloric  acid  (which 
would  result  if  the  proportion  of  potassium  dichromate  to  the  chloride  in  the 
mixture  was  small)  the  reaction  takes  a  different  course,  and  only  chlorine  is 
evolved,  the  chromium  being  completely  reduced  to  chromic  chloride  ;  thus — 

Cr03  +  6HC1  =  3H20  +  CrQ3  +  3C1 
Or,  to  give  the  complete  equation — 

6KC1  +  K2Cr207  +  7H2SO4  -  4K2SO4  -t-  Cr2(SO4)3  +  7H2O  +  3C12 

Hence,  in  applying  this  test,  it  is  necessary  to  employ  an  excess  of  the  dichromate. 
*  The  former  tests  (p.  123),  which  enable  one  to  distinguish  bet-ween  a 
chloride,  bromide,  and  iodide,  will  not  be  confounded  with  a  test  such  as  the 
above,  which  permits  of  the  detection  of  one  class  of  salts  in  the  presence  of 
others. 


Hydrobromic  Acid  and  Bromides.  125 

sulphuric  acid  is  added.  The  beaker  is  then  covered  over  with  a 
piece  of  moistened  blotting-paper  upon  which  a  little  starch  flour 
has  been  dusted.  The  liberated  bromine  produces  a  yellow  colour 
with  the  starch. 

Bromine  may  also  be  liberated  by  means  of  chlorine — 

2KBr  +  C12  =  2KC1  +  Br2 

A  small  quantity  of  carbon  disulphide  is  added  to  the  solution  of 
the  bromide  in  a  narrow  test-tube ;  a  few  drops  of  chlorine  water 
are  added,  and  the  mixture  shaken.  The  liberated  bromine  is 
dissolved  by  the  carbon  disulphide,  giving  a  red  or  brownish 
coloured  liquid  (according  to  the  amount  of  bromine),  which  settles 
to  the  bottom  of  the  test-tube.  The  test  must  be  made  with  a  little 
care,  an  excess  of  chlorine  being  avoided,  as  otherwise  chloride  of 
bromine  is  formed,  which,  being  colourless,  destroys  the  test. 

Hydrobromic  Acid  and  Bromides. 

Gaseous  hydrobromic  acid  closely  resembles  hydrochloric  acid. 
The  properties  of  the  gas  are  not  used  in  analysis. 

The  detection  of  free  hydrobromic  acid  in  solution  in  presence 
of  bromides  may  be  accomplished  by  gently  warming  the  liquid 
with  manganese  dioxide,  without  the  addition  of  sulphuric  acid. 
Bromine  is  liberated  from  the  free  acid  (not  from  the  bromide), 
and  may  be  detected  by  the  starch  reaction  given  above. 

All  bromides  are  soluble  in  water,  except  mercurous  bromide  and 
silver  bromide  ;  lead  bromide  dissolves  in  boiling  water  less  easily 
than  the  chloride. 

Bromides  (except  Hg2Br2  and  AgBr),  when  acted  upon  by 
strong  sulphuric  acid,  evolve  hydrobromic  acid,  bromine,  and 
sulphur  dioxide  (contrast  chlorides,  which  under  these  circum- 
stances give  only  HC1),  the  hydrobromic  acid  first  formed  being 
immediately  acted  upon  by  the  sulphuric  acid  ;  thus — 

KBr  +  H2SO4  =  HBr  +  HKSO4 
2HBr  +  H2SO4  =  Br2  +  SO2  +  2H2O 

Silver  Nitrate,  AgNO3,  precipitates  from  solutions  of 
bromides  or  hydrobromic  acid,  pale-yellow  silver  bromide,  AgBr 
(the  colour  is  indistinguishable  from  white  by  gaslight).  It  is 
insoluble  in  nitric  acid,  and  difficultly  soluble  in  ammonia  (scarcely 
soluble  in  dilute  ammonia.  Contrast  AgCl). 

AgBr  may  be  distinguished  from  AgCl  by  shaking  up  a  little 
of  the  washed  precipitate  with  a  few  drops  of  carbon  disulphide 
and  chlorine  water. 


1 26  Qualitative   A  nalysis. 

Silver  bromide  is  decomposed  by  metallic  zinc  in  the  presence 
of  dilute  sulphuric  acid,  in  the  same  manner  as  the  chloride.  Zinc 
bromide  goes  into  solution,  from  which  the  bromine  can  be  sepa- 
rated by  either  of  the  methods  given  under  bromine. 

Prolonged  boiling  with  a  strong  solution  of  sodium  carbonate 
(or,  better,  heating  the  dry  substances  strongly  in  a  glass  tube) 
decomposes  silver  bromide.  On  filtering  (after  extraction  with 
water  in  the  case  of  the  dry  reaction),  the  aqueous  solution  con- 
taining sodium  bromide  may  be  tested  as  above. 

A  bromide  may  be  detected  in  the  presence  of  a  chloride  by 
means  of  chlorine  water,  which  liberates  bromine  from  the  bromide, 
but  obviously  is  without  action  upon  the  chloride. 

Iodine. 

Iodine  is  a  steel-black,  shining,  crystalline  solid.  When  gently 
heated  it  melts  and  passes  into  vapour,  which  has  a  characteristic 
deep  violet  colour.  Iodine  is  very  slightly  soluble  in  water,  but 
readily  dissolves  in  water  holding  hydriodic  acid  or  alkaline  iodides 
in  solution,  giving  a  brown  liquid.  It  dissolves  in  carbon  disulphide 
(also  in  chloroform),  yielding  a  violet  solution.  In  contact  with 
starch  it  produces  an  intense  indigo-blue  colour,  which  constitutes 
one  of  its  most  delicate  tests. 

Liberation  of  Iodine  from  its  Compounds. — Iodine  is 
more  easily  set  free  from  combination  than  either  bromine  or 
chlorine,  and  the  methods  which  are  applicable  for  the  liberation 
of  these  apply  also  in  the  case  of  iodine.  Thus,  manganese  dioxide 
and  dilute  sulphuric  acid  decompose  iodides  in  a  manner  precisely 
similar  to  that  explained  on  p.  122  for  chlorides.  Strong  acids,  as 
nitric  and  sulphuric,  also  expel  iodine  from  iodides,  with  evolution 
of  oxide  of  nitrogen,  or  sulphur  dioxide  ;  e.g. — 

2K1  +  2H2SO4  =  K2SO4  +  2H2O  +  SO2  +  I2 

The  comparative  ease  with  which  iodine  is  liberated  from  com- 
bination, affords  the  basis  of  most  of  the  tests  by  which  this  element 
is  detected.  The  following  are  the  reactions  most  used  in  analysis : — 

i.  Chlorine  water,  when  added  to  a  solution  of  an  iodide, 
expels  the  iodine.  The  test  may  be  applied  as  described  under 
bromine,  the  carbon  disulphide  in  this  case  being  coloured  violet. 

The  presence  of  the  liberated  iodine  may  also  be  recognised  by 
means  of  starch.  A  small  quantity  of  starch  paste  *  is  mixed  with 

•  *  Starch  paste  is  made  by  mixing  a  little  starch  flour  into  a  thin  cream  with 
the  least  quantity  of  cold  water,  and  then  pouring  boiling  water  upon  it  until  the 
dead-white  appearance  of  the  starch  is  changed  to  a  translucent  appearance. 


Iodine.  127 

the  solution  of  the  iodide,  and  one  or  two  drops  of  chlorine  water 
added,  when  the  deep  indigo-blue  compound  of  iodine  with  starch 
is  produced.  On  the  addition  of  an  excess  of  chlorine,  the  colour 
is  destroyed.  Boiling  the  liquid  also  destroys  the  compound,  hence, 
when  small  quantities  of  iodine  are  being  tested,  it  is  necessary  to 
avoid  using  the  starch  while  hot. 

Detection  of  Bromides  and  Iodides  in  Solution  to- 
gether.— Although'  chlorine  water  liberates  both  bromine  and 
iodine,  the  reaction  may  be  employed  to  detect  both  halogens  in  the 
same  solution,  owing  to  the  fact  that  the  chlorine  exerts  a  selective 
action,  expelling  first  the  iodine,  and  afterwards  the  bromine.  The 
test  is  applied  in  the  following  way  :  Carbon  disulphide  is  added  to 
the  mixed  solution  of  iodide  and  bromide  in  a  test-tube,  and  chlorine 
water  added  in  small  quantities  at  a  time,  with  agitation.  If  this  be 
done  carefully,  it  is  not  difficult  to  see  when  the  further  addition  of 
a  drop  of  chlorine  water  produces  no  further  precipitation  of  iodine. 
At  this  point  the  carbon  disulphide  (coloured  deep  violet  with 
dissolved  iodine)  is  removed  from  the  aqueous  liquor  either  by 
decanting  the  latter,  or,  better,  by  withdrawing  a  portion  of  it  by 
means  of  a  small  pipette  and  transferring  it  to  another  test-tube. 
A  fresh  -quantity  of  carbon  disulphide  is  now  added  to  this,  and  a 
few  drops  of  chlorine  water.  If  the  whole  of  the  iodine  had  been 
liberated  in  the  first  tube,  the  bromine  now  begins  to  be  expelled, 
and  the  carbon  disulphide  becomes  brown.  If  a  small  quantity  of 
iodine  were  still  left,  the  first  drop  of  chlorine  water  causes  its 
liberation,  and,  on  snaking,  the  disulphide  will  show  a  pale  violet 
colour.  A  few  more  drops  of  chlorine  water,  however,  will  destroy 
this,  and  afterwards  liberate  the  bromine.* 

2.  Nitrous  Acid. — When  a  solution  of  an  iodide  is  acidified 
(preferably  with  dilute  sulphuric  acid),  and  a  few  drops  of  a  solution 
of  sodium  nitrite  added,  the  nitrous  acid  generated  (by  the  action 
of  the  acid  upon  the  nitrite)  decomposes  the  iodide,  with  the 
liberation  of  iodine.  .  The  action  is  in  reality  between  the  nitrous 
acid  and  hydriodic  acid  ;  thus — 

HN02  +  HI  =  H20  +  NO  +  I 
or,  to  state  the  complete  change — 
NaNO2  +  2H2SO4  +  KI  =  HNaSO4  +  HKSO4  +  H2O  +  NO  +1 

*  When  the  quantity  of  iodide  present  is  very  small  in  proportion  to  the 
bromide,  it  is  not  necessary  to  change  the  carbon  disulphide  as  here  described. 
The  chlorine,  after  liberating  the  iodine,  next  begins  to  combine  with  it,  form- 
ing iodine  trichloride  (therefore  destroying  the  violet  colour  of  the  carbon 
disulphide  solution),  and  afterwards  to  liberate  the  bromine.  But,  if  the  pro- 
portion of  iodine  present  is  at  all  considerable,  the  yellow  colour  which  the 
iodine  chloride  itself  imparts  to  the  carbon  disulphide — if  present  in  any 
quantity — renders  it  practically  impossible  to  distinguish  the  brownish  colour 
which  would  be  given  by  moderate  quantities  of  bromine. 


128  Qualitative  Analysis. 

Neither  bromine  nor  chlorine  is  liberated  by  nitrous  acid,  hence 
this  reaction  allows  of  the  separation  of  iodine  from  bromides  and 
chlorides,  and  therefore  the  subsequent  detection  of  bromine  or 
chlorine  in  mixtures  of  the  three  salts. 

Detection  of  Iodides,  Bromides,  and  Chlorides  in  Solu- 
tion together. — The  solution  containing  the  three  salts,  to  which 
a  little  carbon  disulphide  has  been  added,  is  acidified  with  two 
or  three  drops  of  dilute  sulphuric  acid,  and  a  dilute  solution  of 
sodium  nitrite  added  drop  by  drop  until  the  whole  of  the  iodine  has 
been  expelled.  On  shaking  the  mixture  this  will  be  dissolved  by  the 
carbon  disulphide,  giving  the  violet  solution.  The  aqueous  liquid 
is  then  withdrawn  with  a  pipette  and  divided  into  two  portions. 
The  first  is  neutralised  by  the  cautious  addition  of  ammonia  drop 
by  drop.  It  is  then  shaken  with  chlorine  water  and  carbon  disul- 
phide. The  bromine  is  thereby  liberated,  and  imparts  its  brownish 
colour  to  the  disulphide.  The  second  portion  is  evaporated  down, 
mixed  with  potassium  dichromate  and  sulphuric  acid,  and  the 
chromyl-chloride  test  made  as  described  on  p.  123. 

3.  Ferric  Chloride. — When  ferric  chloride,  acid  with  hydro- 
chloric acid,  is  added  to  a  solution  of  an  iodide  and  the  -mixture 
heated,  the  iron  salt  is  reduced  to  the  ferrous  condition,  and  iodine 
escapes  as  a  violet  vapour — 

FeCl3  +  KI  =  KC1  +  FeCl2  +  I 

Hydriodic  Acid  and  Iodides. 

Gaseous  hydriodic  acid  resembles  hydrobromic  and  hydro- 
chloric acids,  except  that  it  is  more  easily  decomposed,  being 
dissociated  into  iodine  and  hydrogen  by  very  moderate  heat.  The 
properties  of  the  gas  are  not  used  in  analysis. 

All  iodides  are  soluble  in  water,  except  those  of  silver,  Agl  ; 
mercury,  Hg2I2  and  HgI2 ;  copper,  Cu2I2  ;  gold,  AuI3 ;  platinum, 
PtI4  ;  palladium,  PdI2;  (BiI2  and  PbI2  sparingly  soluble). 

The  formation  of  some  of  these  insoluble  salts  is  utilised  in 
analysis. 

Silver  nitrate,  AgN03,  when  added  to  a  solution  of  hydriodic 
acid  or  an  iodide,  gives  a  pale-yellow  precipitate  of  silver  iodide, 
Agl,  insoluble  in  nitric  acid,  and  more  difficult  of  solution  in 
ammonia  than  silver  bromide.  [In  very  strong  solutions  the  pre- 
cipitation is  incomplete,  owing  to  the  partial  solubility  of  Agl  in 
concentrated  solutions  of  HI  or  KI.  Dilution  with  water  repre- 
cipitates  the  dissolved  compound.] 

Silver  iodide  may  be  distinguished  from  the  bromide  or  chloride 


Hypochlorous  Acid.  129 

by  shaking  up  the  precipitate  with  a  little  carbon  disulphide  and 
chlorine  water. 

Silver  iodide  is  reduced  by  metallic  zinc  in  the  same  way  as  the 
bromide  and  chloride,  or  it  may  be  decomposed  (in  common  with 
all  the  other  insoluble  iodides)  by  fusion  with  sodium  carbonate. 

Copper  sulphate,  CuSO4,  added  to  a  solution  of  an  iodide, 
gives  a  dirty  brown  precipitate,  consisting  of  cuprous  iodide,  Cu2I2, 
and  free  iodine  (see  Copper  reactions,  p.  87) — • 

2CuSO4  +  4KI  =  2K2SO4  +  Cu2I2  +  I2 
Cuprous  iodide  is  a  white  crystalline  compound,  insoluble  in 

water  and  in  dilute  acids  ;  soluble  in  ammonia  and  in  sodium  thio- 

sulphate. 

In  the  presence  of  suitable  reducing  agents,  the  whole  of  the 

iodine  may  be  precipitated  as  cuprous  iodide  ;  and  as  bromine  and 

chlorine  are  not  precipitated  under  these  conditions,  this  reaction 

is  made  use  of  as  a  means  of  separation. 

Separation  of  Iodine  as  Cuprous  Iodide.— To  the  solution 
containing  an  iodide,  bromide,  and  chloride,  copper  sulphate  is  added 
in  excess  (i.e.  in  quantity  more  than  sufficient  to  combine  with  the 
whole  of  the  iodine).  Sodium  thiosulph«..e  is  added  drop  by  drop 
from  a  pipette  until  the  brown  colour  (due  to  the  free  iodine) 
disappears,  and  the  precipitate  is  white.  When  this  is  filtered, 
the  liquid  will  be  coloured  blue  with  the  excess  of  copper  sulphate, 
and  will  contain  the  bromide  and  chloride,  for  which  it  may  be 
examined  by  methods  already  described. 

The  first  action  of  the  sodium  thiosulphate  is  to  convert  the 
free  iodine  into  sodium  iodide,  and  to  form  sodium  tetra-thionate — 

2Na2S2O3  -f  I2  =  2NaI  +  Na2S4O0 

and  the  sodium  iodide,  then  reacting  upon  copper  sulphate  in  the 
presence  of  the  thiosulphate,  has  the  whole  of  its  iodine  precipi- 
tated as  Cu2I2. 

Fallacious  nitrate,  Pd(NO3)2,  when  added  to  a  solution 
of  an  iodide,  gives  a  black  precipitate  of  palladious  iodide,  PdI2. 
The  reaction  is  delicate,  notwithstanding  that  the  precipitate  is 
soluble  to  some  extent  in  an  excess  of  potassium  iodide.  It  serves 
to  detect  iodides  in  the  presence  of  chlorides  and  bromides. 

THE  HALOGEN  OXYACIDS. 
Hypochlorous  Acid  and  Hypochlorites. 

Hypochlorous  acid,  HC1O,  is  only  known  in  solution  in 
water,  as  it  readily  undergoes  decomposition.  The  solution  has  a 
pale  yellow  colour,  and  a  faint  smell  (familiar  in  ordinary  chloride 
of  lime].  It  also  has  bleaching  properties. 

K 


130  Qualitative  Analysis. 

A  solution  of  the  free  acid  may  be  distinguished  from  chlorine 
water  (besides  from  its  odour)  by  shaking  it  up  with  a  little  mer- 
cury. With  chlorine  water  white  mercurous  chloride,  Hg2Cl2,  is 
formed,  while  with  hypochlorous  acid  a  brownish-coloured  oxy- 
chloride  is  produced — 

2Hg  +  2HC10  -  HgO,HgCl2  +  H20 

All  hypochlorites  are  soluble  in  water,  therefore  no  precipitation 
reactions  for  their  detection  are  possible.  The  tests  employed  for 
their  identification  are  based  upon  their  oxidising  action.  They 
are  decomposed  by  even  feeble  acids  (carbonic),  with  evolution  of 
chlorine,  e.g. — 

NaCIO  +  2HC1  =  NaCl  +  H2O  +  C12 
Ca(ClO)2  4-  CO2  =  CaCO3  +  C12  +  O 

If  a  few  drops  of  a  solution  of  a  hypochlorite  be  poured  upon 
litmus  paper,  and  the  paper  exposed  to  the  action  of  carbon  dioxide 
(by  breathing  upon  it),  it  will  become  bleached  where  it  was  mois- 
tened, by  the  evolved  chlorine. 

The  evolution  of  chlorine  by  the  action  of  dilute  acids  distin- 
guishes hypochlorites  from  chlorides  and  chlorates.  The  addition 
of  silver  nitrate  to  a  hypochlorite  results  in  the  precipitation  of 
silver  chloride,  one-third  of  the  silver  going  into  solution  as  silver 
chlorate  ;  thus — 

(1)  NaCIO  +  AgNO3  =  AgCIO  +  NaNO3 

(2)  3AgC10  =  AgClOs  +  2AgCl 

The  commonest  salt  of  hypochlorous  acid  is  bleaching  powder, 
which  is  a  double  chloride  and  hypochlorite  of  calcium,  having  the 
composition  Ca(OCl)Cl.  When  treated  with  water  it  splits  up 
into  calcium  chloride  and  calcium  hypochlorite,  and  a  clear  solu- 
tion made  in  this  way  contains  the  salt  Ca(ClO)2 — 
2Ca(ClO)Cl  =  CaCl2  +  Ca(ClO)2 

Chlorates,  Bromates,  and  lodates. 

The  three  acids,  chloric,  HC1O3,  bromic,  HBrO3,  and  iodic, 
HIO3,  differ  in  stability  in  the  opposite  order  to  that  usually 
exhibited  by  compounds  of  the  halogens  ;  thus,  iodic  acid  is  a 
comparatively  stable  solid,  while  chloric  acid  can  only  exist  in 
aqueous  solutions  containing  not  more  than  20  per  cent,  of  the 
acid.  Beyond  this  strength,  or  when  heated,  it  decomposes  into 
oxygen,  chlorine,  and  perchloric  acid — 

3HC103  =  2O2  +  C12  +  HC104  +  H2O 


Chlorates,  Bromates,  and  lodates.  131 

Bromic  acid  under  the  same  conditions  breaks  up  into  bromine, 
oxygen,  and  water  ;  thus  — 

4HBrO3  =  5O2  +  2Br2  +  2H2O 

lodic  acid  does  not  give  a  blue  colour  with  starch,  but  it  readily 
parts  with  its  oxygen  to   such  reducing  agents   as   sulphuretted 
hydrogen  or  sulphur  dioxide,  with  liberation  of  iodine  ;  thus  — 
2HIO3  +  5H2S  =  58  +6H2O  +  I2* 

Similarly,   hydriodic    acid   and   iodic   acid    undergo    interaction, 
the  whole  of  the  iodine  being  thrown  down  — 


This  reaction  affords  the  means  of  detecting  the  presence  of  an 
iodate  when  mixed  with  an  iodide.  A  little  starch  is  added,  and  a 
small  quantity  of  acetic  acid  (this,  while  liberating  the  respective 
iodine  acids  from  their  salts,  is  not  able  to  decompose  the  iodic 
acid),  when  the  blue  starch  coloration  is  obtained. 

The  salts  of  these  acids  are  produced  (mixed  with  the  corre- 
sponding halides)  by  the  direct  action  of  the  halogens  upon  caustic 
alkalies  ;  thus,  in  the  case  of  iodine  (bromine  and  chlorine  being 
similar)  — 

6KHO  4-  3l2  =  5KI  +  KIO3  +  3H2O 

Chlorates  are  all  soluble  in  water. 

Bromates.  —  The  silver  and  mercurous  salts  are  difficult  of 
solution,  and  may  be  obtained  by  precipitation. 

lodates.  —  Only  the  iodates  of  the  alkali  metals  are  soluble  in 
water. 

The  chlorates,  bromates,  and  iodates  are  all  decomposed  by 
heat,  yielding  either  oxygen  and  the  haloid  salt  (or  oxygen,  metallic 
oxide,  and  free  halogen)  ;  or,  in  some  cases,  leaving  both  an  oxide 
and  halide  in  the  residue  —  for  example  — 

(1)  AgC103  =  AgCl  +  30 

(2)  Mg(BrO3)2  =  MgO  +  Br2  +  50 

(3)  2Pb(Br03)2  -  PbO  +  PbBr2  +  Br2  +  iiO 

When  heated  with  oxidisable  substances  (e.g.  charcoal),  deflagration 
of  the  mixture  results. 

Hydrochloric  acid  decomposes  chlorates,  with  the  evolution 
of  chlorine  and  chlorine  peroxide  — 

4KC1O3  +  I2HC1  =  4KC1  +  6H2O  +  QC1  +  3C1O2 

*  Owing  to  this  property,  if  iodic  acid  is  present  in  a  mixture  undergoing 
analysis,  iodine  will  be  precipitated  in  the  process  of  separating  Group  II.  by 
means  of  H2S.  By  the  continued  action  of  the  sulphuretted  hydrogen,  however, 
the  iodine  is  converted  into  hydriodic  acid,  H2S  +  I2  =  S  +  2HI. 


132  Qualitative  Analysis. 

(The  use  of  this  mixture  as  an  oxidising  agent  has  frequently  been 
referred  to.) 

Bromates  and  iodates  under  the  same  treatment  give  bromine 
and  iodine  respectively.  In  the  latter  case,  chloride  of  iodine  is 
also  formed  — 

2KBrO3  +  2HC1  =  2KC1  +  H2O  +  Br2  +  50 
KI03  +  6HC1  =  KC1  +  IC13  +  C12+  3H20 

Sulphuric  acid  decomposes  chlorates,  with  the  evolution  of 
chlorine  peroxide  (a  deep-yellow  unpleasant-smelling  gas),  which 
on  very  slight  elevation  of  temperature,  explodes  with  violence  — 

3KC103  +  2H2S04  =  KC104  +  2HKS04  +  H2O  +  2C1O2 

A  minute  crystal  of  potassium  chlorate,  with  about  three  or 
four  drops  of  strong  sulphuric  acid,  may  be  heated  in  a  test-tube  ; 
sharp  detonations  will  result,  characteristic  of  chlorates.  From  bro- 
mates  sulphuric  acid  liberates  bromine.  Dilute  acid  first  liberates 
bromic  acid,  which  gradually  decomposes  into  HBr  and  O.  As 
soon  as  this  decomposition  sets  in,  the  HBr  reacts  upon  the  HBrO3, 
with  liberation  of  bromine  — 


HBr03  +  5HBr  =  3H2O  +  3Br2 

Dilute  sulphuric  acid,  therefore,  liberates  bromine  immediately 
from  a  mixture  of  a  bromide  and  bromate  ;  but  only  gradually,  after 
an  interval,  from  a  bromate  alone. 

Silver  nitrate  gives  with  bromates  a  white  precipitate  of 
silver  bromate,  AgBrO3  ;  soluble  in  ammonia.  It  is  decomposed  by 
hydrochloric  acid,  with  liberation  of  bromine  (distinction  from  all 
silver  halides}.  With  iodates,  silver  nitrate  precipitates  white  silver 
iodate,  AgIO3.  Difficultly  soluble  in  nitric  acid,  readily  dissolved 
by  ammonia  (distinction  from  Agl).  With  chlorates  silver  nitrate 
gives  no  precipitate.* 

Hydrogen  sodium  sulphite  added  to  a  solution  of  an  iodate, 
liberates  the  whole  of  the  iodine  ;  thus  — 

2NaIO3  +  sHNaSO3  =  3HNaSO4  +  2Na2SO4  +  H2O  +  I2 

Similarly,  if  sulphurous  acid  be  added  to  the  ammoniacal  solution 
of  silver  iodate,  the  iodate  is  entirely  reduced,  and  the  iodine 
thrown  out  as  silver  iodide  — 

AgI03  +  3H(NH4)S03  =  3H(NH4)SO4  4-  Agl 

*  For  method  of  detecting  a  chlorate  and  nitrate  in  solution  together,  see 
Nitric  acid  (p.  147). 


Hydrofluoric  Acid  and  Fluorides.  133 

Barium  chloride  precipitates  from  soluble  iodates,  white 
barium  iodate,  Ba(IO3)2.  Slightly  soluble  in  water,  insoluble  on 
addition  of  alcohol ;  slightly  soluble  in  nitric  acid,  easily  soluble 
in  hydrochloric  acid.  (This  reaction  affords  a  means  of  separating 
iodic  from  hydriodic  acid,  as  well  as  from  all  other  halogen  acids.) 

A  chloride  may  be  separated  from  a  chlorate  by  adding  to  the 
solution  silver  sulphate.  Silver  chloride  is  precipitated,  and  removed 
by  filtration.  Sodium  carbonate  is  then  added  to  remove  the 
excess  of  silver,  and  any  metals  other  than  the  alkalies  which 
might  be  present  ;  and  the  solution  evaporated  to  dryness,  and 
heated.  The  chlorate  is  thereby  converted  into  chloride. 

Per  chlorates. 

The  perchlorates  are  all  soluble  in  water,  and  are  more  stable 
than  the  chlorates.  Thus,  when  potassium  chlorate  is  heated,  the 
final  products  of  the  decomposition  are  KC1  and  O.  But  as  an 
intermediate  product  potassium  perchlorate  is  formed  ;  thus — 

2KC1O3  =  KC1O4  +  KC1  +  O2 

and  on  applying  a  higher  temperature,  the  potassium  perchlorate  is 
finally  resolved  into  potassium  chloride  and  oxygen.  (Bromates, 
when  heated,  pass  at  once  into  bromides,  with  evolution  of  oxygen.) 

Sulphuric  acid  liberates  perchloric  acid,  HC1O4,  a  colourless, 
fuming,  corrosive,  volatile  liquid.  The  mixture  of  the  perchlorate 
and  acid  does  not  become  yellow  ;  and  on  warming,  the  acid 
distils  without  decomposition  (distinction  from  chlorates,  also  from 
bromates]. 

Hydrochloric  acid  has  no  action  on  perchlorates  (distinction 
from  chlorates,  bromates,  iodates). 

Fluorine,  Hydrofluoric  Acid,  and  Fluorides. 

The  properties  of  the  element  are  not  made  use  of  in  analysis, 
as  fluorine  is  not  liberated  by  any  of  the  ordinary  analytical  re- 
actions. 

Hydrofluoric  acid  is  a  colourless,  fuming,  highly  corrosive  liquid, 
boiling  at  19-5°  C,  therefore  extremely  volatile.  It  dissolves  in 
water,  and  even  a  moderately  dilute  solution  attacks  the  skin 
violently.  The  most  characteristic  property  of  the  acid,  either 
gaseous  or  in  solution,  is  its  power  of  dissolving  silica  (therefore 
glass)  :  its  application  is  described  below. 

Liberation  of  Hydrofluoric  Acid. — Fluorides  are  decom- 
posed by  strong  sulphuric  acid,  with  evolution  of  the  acid  in  the 


134  Qualitative  Analysis. 

gaseous  condition.     Thus  with  the  mineral  cryolite^  and  the  still 
commoner  fluor  spar — 

2Na3AlF6  +  6H2SO4  =  3Na2SO4  4-  A12(SO4)3  +  I2HF 
CaF2  +  H2SO4  =  CaSO4  +  2HF 

Hydrofluoric  acid  is  also  evolved  when  a  fluoride  is  heated  with 
powdered  hydrogen  potassium  sulphate — 

CaF2  +  2HKSO4  =  CaSO4  +  K2SO4  +  2HF 

Certain  acid  fluorides,  when  heated  alone,  decompose  with 
evolution  of  hydrofluoric  acid  ;  thus — 

HF,KF  =  KF  +  HF 

The  fluorides  of  the  non-metals  are  all  volatile  without  decom- 
position. 

The  fluorides  of  the  alkali  metals,  and  of  silver,  mercury  (iron, 
aluminium,  tin),  are  soluble  in  water.  Those  of  the  alkaline  earths, 
and  of  lead  (copper,  zinc,  manganese),  are  insoluble. 

Calcium  chloride  gives  a  transparent  gelatinous  precipitate 
of  calcium  fluoride,  CaF2,  partially  soluble  in  hydrochloric  acid. 

Barium  chloride  throws  down  a  white  precipitate  of  barium 
fluoride,  BaF2,  partially  soluble  in  HC1. 

Silver  nitrate  gives  no  precipitate,  as  silver  fluoride  is  soluble 
in  water  (distinction  between  a  fluoride  and  the  other  halides). 

Formation  of  Silicon  Fluoride. — When  hydrofluoric  acid 
comes  in  contact  with  silica,  or  silicates  (such  as  glass),  the  silicon 
dioxide  is  dissolved,  and  gaseous  silicon  fluoride  is  formed — 

SiO2  +  4HF  =  SiF4  +  2H2O 

The  test  may  be  applied  in  the  following  ways  : — 
(a)  Etching  Glass. — The  powdered  fluoride  is  mixed  with 
strong  sulphuric  acid  in  a  small  dish  or  tray,  made  of  lead  (or  a 
platinum  capsule).  It  is  covered  with  a  small  piece  of  sheet  glass 
which  has  been  coated  on  one  side  with  wax,*  and  some  marks  or 
words  scratched  upon  the  wax.  In  a  few  minutes  the  exposed 
parts  of  the  glass  will  have  become  eaten  into  or  dissolved  away 
by  the  acid  gas  ;  so  that,  on  removing  the  wax  with  a  little  hot 
water,  the  marks  or  letters  will  be  found  to  be  etched  into  the  glass. 
(£)  The  Decomposition  of  Silicon  Fluoride  by  Water.— 
If  the  fluoride  (natural  fluoride)  contains  much  silica,  or  if  it  be 

*  A  little  paraffin  wax  is  melted  in  a  test-tube,  and  poured  upon  the 
previously  warmed  glass  plate.  The  liquid  wax  is  made  to  flow  all  over  the 
surface  of  the  plate,  which  is  then  stood  up  on  edge  to  drain  and  cool.  In  this 
way  a  thin  and  uniform  film  of  wax  will  be  obtained. 


Hydrofluoric  Acid.  135 

intentionally  mixed  with  silica  (sand),  the  action  of  sulphuric  acid  is 
to  liberate  silicon  fluoride,  a  gas  which  is  without  action  upon  glass. 
The  hydrofluoric  acid  (from  the  fluoride),  being  generated  in  the 
immediate  presence  of  silica,  at  once  combines  with  it ;  thus — 
2CaF2  +  2H2SO4  +  SiO2  =  2CaSO4  +  2H2O*  +SiF4 

This  gas  is  at  once  decomposed  by  water,  with  the  formation  of 
hydrofluosilicic  acid  (soluble  in  water)  and  the  precipitation  of 
gelatinous  silicic  acid,  H2SiO3 — 

3SiF4  +  3H20  =  2H2SiF6  +  H2Si03 

The  mixture  of  the  fluoride  and  sand  is  gently  warmed  in  a 
test-tube  with  a  little  strong  sulphuric  acid,  and  a  glass  rod  with  a 
drop  of  water  upon  the  end  is  lowered  into  the  mouth  of  the  tube. 
The  gas,  on  coming  in  contact  with  the  water,  is  decomposed,  and  a 
v/hite  deposit  of  silicic  acid  is  formed  upon  the  rod.  [The  experi- 
ment is  rendered  more  delicate  by  using  a  piece  of  narrow  glass 
tube  with  a  drop  of  water  at  the  end.  By  means  of  a  piece  of 
rubber  pipe  attached  to  the  other  end  of  the  tube,  the  drop  of  water 
can  be  slowly  sucked  up  the  tube,  and  the  gas  which  is  drawn  up 
after  it  will  then  deposit  a  white  film  of  silica  upon  the  wet  walls  of 
the  tube.  The  gas  must  not  be  drawn  into  the  lungs.]  (See  also 
Hydrofluosilicic  acid.) 

The  Formation  of  Boron  Fluoride.— When  a  fluoride  (finely 
powdered)  is  mixed  with  powdered  borax,  and  the  mixture  moistened 
with  strong  sulphuric  acid,  gaseous  boron  fluoride,  BF3,  is  evolved. 
The  action  takes  place  between  the  hydrofluoric  acid  and  boric 
acid,  which  are  disengaged  by  the  action  of  the  sulphuric  acid  upon 
the  respective  compounds — 

B(HO)3  +  3HF  =  3H20  +  BF3 

If  the  mixture  be  introduced  into  the  edge  of  a  Bunsen  flame 
upon  a  loop  of  platinum  wire,  the  flame  is  tinged  a  grass-green  colour 
by  the  escaping  boron  fluoride. 

[The  test  may  be  modified  by  employing  hydrogen  potassium 
sulphate  instead  of  sulphuric  acid,  and  heating  the  mixture  of  the 
three  salts  upon  the  platinum  wire.] 

Hydrofluosilicic  Acid  and  Silicofluorides. 

This  acid  is  obtained  when  silicon  fluoride  is  decomposed  by 
water.  The  gas  is  made  to  bubble  into  water,  and  the  silicic  acid 
separated  by  filtration. 

*  The  two  molecules  of  H2O  here  produced  are  not  liberated  as  water,  but 
are  retained  by  the  calcium  sulphate  in  the  form  of  water  of  hydration  ;  the 
fully  hydrated  compound  having  the  composition  CaSO4,2H2O. 


136  Qualitative  Analysis. 

The  acid  and  its  salts  are  decomposed  by  heat-^ 

H2SiF6  =  SiF4  +  2HF 
BaSiF6  =  SiF4  +  BaF2 

The  silicofluorides  of  potassium,  sodium,  and  barium  are  pre- 
cipitable  salts,  being  only  difficultly  soluble  in  water.  The  addition 
of  alcohol  renders  the  precipitation  complete  (see  Reactions  of 
metals  of  Group  IV.  and  V.).  Barium  silicofluoride  is  only 
sparingly  soluble  in  dilute  hydrochloric  acid,  and  is  therefore  pre- 
cipitated by  barium  chloride  in  a  hydrochloric  acid  solution. 
Barium  silicofluoride,  however,  is  distinguished  from  barium  sulphate 
by  dissolving  in  strong  hydrochloric  acid  (the  solution  being  after- 
wards diluted  to  dissolve  any  barium  chloride  which  may  be  thrown 
down  by  the  strong  acid)  ;  or  by  the  fact  that  when  dried  and 
heated  it  evolves  silicon  fluoride,  as  shown  by  the  above  reaction. 

The  addition  of  ammonia. to  hydrofluosilicic  acid,  or  a  soluble 
silicofluoride,  gives  a  white  precipitate  consisting  of  gelatinous 
silicic  acid — 

K2SiF6  +  4NH4HO  =  2KF  +  4NH4F  +  H2O  +  H2SiO3 


CHAPTER  XII. 

SULPHUR. 

THE  chief  properties  made  use  of  in  analysis  are  the  following  : 
Sulphur  is  a  pale-yellow,  brittle  solid.  Insoluble  in  water  ;  easily 
soluble  in  carbon  disulphide,  from  which  it  is  deposited  on  evapora- 
tion in  amber-coloured  rhombic  octahedrons. 

When  heated  in  a  test-tube,  it  melts  and  boils  off  as  a  brownish- 
yellow  vapour,  which  condenses  to  brown  drops  on  the  moderately 
hot  parts  of  the  tube  (turning  yellow  when  cold),  and  as  a  yellow 
sublimate  on  the  upper  and  cooler  part  of  the  tube  (flowers  of 
sulphur). 

When  heated  with  access  of  air,  either  on  a  platinum  capsule  or 
in  a  glass  tube  open  at  both  ends,  sulphur  burns  with  a  pale  blue 
flame,  producing  sulphur  dioxide  (of  characteristic  smell),  and 
leaving  no  residue.  Nitric  acid  oxidises  sulphur  into  sulphuric 
acid,  with  evolution  of  nitrogen  peroxide. 

Liberation  of  sulphur  takes  place  when  certain  metallic 
sulphides  are  heated  alone,  either  in  a  tube  closed  at  one  end,  or 
with  partial  access  of  air  ;  e.g. — 

3FeS2  =.-.  Fe3S4  +  S2 
3FeS2  +  5O2  -  Fe3O4  +  380,  +  38 

Sulphuretted  Hydrogen*  and  Sulphides. 

Sulphuretted  hydrogen  is  a  colourless  gas,  easily  distinguished 
from  all  other  gases  by  its  unmistakable  odour.  It  is  soluble  in 
water,  and  imparts  its  own  smell  to  the  liquid.  The  solution,  how- 
ever, is  unstable,  undergoing  oxidation  and  depositing  sulphur. 
The  gas  burns  with  a  flame  resembling  that  of  burning  sulphur, 
and  yields  water  and  sulphur  dioxide. 

Liberation  of  sulphuretted  hydrogen  takes  place  when 

*  Sometimes  called  hydrosulphuric  acid ;  the  solution  of  the  gas  in  water 
has  a  feeble  acid  reaction.  It  should  be  remembered  that  the  gas  is  poisonous. 


138  Qualitative  Analysis. 

certain  sulphides  (see  below)  are  acted  upon  by  acids.  The  gas  may 
be  recognised  (i)  by  its  odour  ;  (2)  by  its  action  upon  solutions  of 
metallic  salts,  e.g.  lead  acetate.  The  reaction  is  made  in  a  test- 
tube,  and  a  piece  of  paper  moistened  with  lead  acetate  is  held  over 
the  mouth  of  the  tube.  The  sulphuretted  hydrogen  causes  a  black 
stain  of  lead  sulphide. 

Sulphides  of  the  alkalies  and  alkaline  earths  are  soluble  in 
water ;  all  other  metallic  sulphides  are  insoluble  (see  Analytical 
classification  of  the  metals). 

Soluble  sulphides  are  decomposed  by  dilute  acids  (HC1  or 
H2SO4),  with  liberation  of  sulphuretted  hydrogen;  in  the  case  of 
polysulphides,  sulphur  is  also  precipitated.* 

K2S  +  2HC1  +  Aq  =  2KC1  +  H2S  +  Aq 
CaS5  -t-  2HC1  =  CaCl2  +  H2S  +48 

Insoluble  Sulphides.!— The  behaviour  of  these  towards 
acids  has  already  been  considered  in  detail,  in  studying  the  separa- 
tion of  the  metals.  It  may  be  briefly  summarised  as  follows  : — 

(a)  Sulphides  decomposed  by  dilute  acids  (HC1  or  H2SO4),  with 
liberation  of  sulphuretted  hydrogen  :  namely,  ZnS,  MnS,  FeS. 

(b)  Sulphides  unacted  upon  by  dilute  acid,  but  decomposed  by 
hot  strong  hydrochloric  acid  with  more  or  less  difficulty :  Sb2S3, 
PbS,  SnS,  NiS,  CoS. 

(c)  Sulphides  unacted  upon    by  strong  hydrochloric  acid,  but 
decomposed  by  aqua  regia,  or  by  a  mixture  of  hydrochloric  acid 
and  potassium  chlorate  ;  HgS,  As2S3,  PtS2,  AuS. 

The  sulphides  of  class  (b},  when  treated  with  hydrochloric  acid 
in  the  presence  of  zinc,  or,  better,  of  reduced  iron,  readily  evolve 
sulphuretted  hydrogen. 

Oxidising  agents,  e.g.  nitric  acid,  convert  many  of  the  sul- 
phides into  oxides  or  sulphates,  sulphur  being  first  separated  and 
afterwards  oxidised  into  sulphuric  acid. 

When  a  sulphide  is  added  (in  small  quantities  at  a  time)  to  a 
fused  mixture  of  sodium  carbonate  and  potassium  nitrate  in  a 
platinum  crucible,  the  sulphide  is  immediately  oxidised.  After  the 
mass  has  cooled,  and  been  extracted  with  water,  the  aqueous  liquid 
may  be  tested  for  a  sulphate.  The  sulphides  of  classes  (a)  and  (£), 
when  fused  upon  a  piece  of  platinum  foil  (or,  better,  silver)  with 

*  Under  certain  conditions  hydrogen  persulphide  is  formed,  without  any 
evolution  of  sulphuretted  hydrogen. 

f  Aluminium  sulphide,  A12S3,  is  not  formed  in  analysis.     It  is  decomposed 
by  water,  with  the  liberation  of  sulphuretted  hydrogen  ;  thus— 
A12S3  +  sH20  =  A1203  +  3H2S 


Stilphuric  Acid  and  Sulphates.  139 

sodium  hydroxide,  are  decomposed,  with  the  formation  of  sodium 
sulphide.  If  a  fragment  of  the  fused  mass,  after  cooling,  be  placed 
upon  a  silver  coin  and  moistened  with  a  drop  of  water,  or  upon 
a  piece  of  paper  which  has  been  moistened  with  a  solution  of  lead 
acetate,  in  either  case  a  black  stain  will  be  produced  ;  silver  sulphide 
on  the  coin,  and  lead  sulphide  upon  the  paper. 

Most  sulphides,  when  heated  in  a  glass  tube  open  at  both  ends, 
and  held  in  a  slightly  inclined  position  in  order  to  cause  an  air- 
current  to  pass  through  the  tube,  are  decomposed,  and  evolve  sul- 
phur dioxide. 

Sodium  nitroprusside,  Na2(NO)FeCy5)*  gives  with  alkaline 
sulphides  (but  not  with  an  aqueous  solution  of  sulphuretted  hydro- 
gen] a  rich  reddish-purple  colour.  One  or  two  drops  of  the  nitro- 
prusside solution  are  added  to  the  alkaline  sulphide  (e.g.  ammonium 
sulphide)  upon  a  white  plate.  The  colour  is  destroyed  by  caustic 
alkalies,  which  must  therefore  be  absent  when  applying  the  test. 
(By  this  test,  a  soluble  sulphide  may  be  detected  in  the  presence  of 
dissolved  sulphuretted  hydrogen.) 

Sulphuric  Acid  and  Sulphates. 

Sulphuric  acid  is  an  oily,  highly  corrosive  acid  liquid.  It  com- 
bines with  water  with  evolution  of  heat,  and  is  able  to  abstract  the 
elements  of  water  from  many  organic  compounds.  Thus  paper, 
straw,  etc.,  are  blackened  or  charred  by  the  strong  acid.  This  pro- 
perty is  made  use  of  in  testing  for  the  free  acid  in  the  presence  of 
soluble  sulphates  :  a  piece  of  paper  is  moistened  here  and  there 
with  the  solution,  and  then  carefully  dried,  when  it  becomes  charred 
where  it  had  been  wetted.  Or  the  solution  may  be  mixed  with  a 
little  white  sugar,  and  evaporated  down  in  a  porcelain  dish  upon  a 
steam-bath,  when  a  charred  residue  will  be  left. 

Most  sulphates  are  soluble  in  water.  Barium,  strontium,  cal- 
cium, and  lead  sulphates  are  insoluble,  or  nearly  so  (see  Reactions 
of  the  various  metals).  All  sulphates  (except  ferric  sulphate)  are 
precipitated  by  alcohol,  being  insoluble  in  that  liquid  (hence  the 
use  of  alcohol  in  rendering  the  precipitation  of  the  sulphates  of 
lead,  calcium,  or  strontium  complete). 

Soluble  Sulphates.— Barium  chloride,  BaCl2,  gives  with  sul- 
phuric acid  or  soluble  sulphates,  a  white  precipitate  of  barium 

*  This  compound  is  readily  obtained  by  boiling  a  little  solid  potassium 
ferrocyanide  with  strong  nitric  acid,  diluting  with  water,  and  neutralising  with 
sodium  carbonate.  The  salt  easily  crystallises  in  deep  red  crystals.  The  exact 
nature  of  the  action  of  alkaline  sulphides  upon  it  is  unknown. 


140  Qualitative  Analysts. 

sulphate  (see  Barium  reactions),  insoluble  in  hydrochloric  acid.* 
The  solutions  should  be  dilute,  as  barium  chloride,  being  insoluble 
in  strong  hydrochloric  acid,  may  otherwise  be  thrown  out  of  solu- 
tion ;  the  addition  of  water  dissolves  it. 

Insoluble  sulphates  may  be  decomposed  by  fusion  with 
sodium  carbonate,  sodium  sulphate  being  formed.  The  residue  is 
extracted  with  water,  and  the  aqueous  solution  tested  with  barium 
chloride  after  being  acidified. 

When  a  sulphate  is  fused  with  sodium  carbonate  (which  must 
be  free  from  sulphates  as  impurities)  upon  charcoal  in  the  reducing 
flame,  a  sulphide  of  the  alkali  metal  is  obtained.  If  this  be  placed 
upon  a  piece  of  paper  moistened  with  acetate  of  lead,  and  touched 
with  a  drop  of  dilute  hydrochloric  acid,  sulphuretted  hydrogen  is 
liberated,  and  the  lead  paper  stained  black. 

[This  test  is  only  conclusive  evidence  of  a  sulphate  when  other 
sulphur  compounds  are  proved  to  be  absent.] 

Sulphurous  Acids  and  Sulphites. 

Sulphurous  acid,  H2SO3,  is  only  known  in  solution,  being  pro- 
duced when  sulphur  dioxide  is  passed  into  water,  or  when  this  gas 
is  liberated  from  combination  (as  from  sulphites)  in  the  presence 
of  water  ;  thus — 

(1)  In  dilute  solution,  Na2SO3  + 2HC1  +Aq  =  2NaCl+H2SO3  +  Aq 

(2)  In  stronger  solution,  Na2SO3  +  2HC1  =  2NaCl  +H2O  +  SO2 

The  anhydride,  SO2,  is  recognised  by  its  characteristic  suffocating 
odour  (familiar  as  "  the  smell  of  burning  sulphur  "). 

Reducing  Action  of  Sulphurous  Acid. — Sulphurous  acid 
easily  takes  up  oxygen,  and  passes  into  sulphuric  acid,  and  some 
of  its  most  important  reactions  are  those  in  which  it  thus  acts  as  a 
reducing  agent.  Many  of  these  have  been  referred  to  under  the 
metals  ;  thus,  potassium  permanganate  is  reduced  with  formation 
of  manganous  sulphate — 

2KMnO4  +  5H2SO3  =  K2SO4  +  2MnSO4  +  2H2SO4  +  3H2O 

This  reaction  affords  a  delicate  test  for  sulphur  dioxide.  The  gas 
is  cautiously  decanted  (being  much  heavier  than  air)  into  a  test- 
tube  containing  water  slightly  tinted  with  a  minute  quantity  of 
potassium  permanganate.  On  shaking  the  gas  and  water,  the  pink 
colour  will  be  destroyed. 

*  The  insolubility  of  barium  sulphate  in  hydrochloric  acid  does  not  dis- 
tinguish this  compound  from  barium  silicofluoride.  The  latter  compound, 
however,  when  dried  and  heated,  gives  off  silicon-fluoride  (see  p.  136). 


Sulphurous  Acid  and  Sulphites.  141 

Oxidising  Action  of  Sulphurous  Acid.—  Sulphurous  acid  is 
also  capable  of  undergoing  reduction,  acting  therefore  towards  more 
powerful  reducing  agents  in  the  capacity  of  an  oxidising  substance. 
Thus,  stannous  chloride  in  presence  of  hydrochloric  acid,  is  oxidised 
into  stannic  chloride,  the  sulphurous  acid  being  reduced  to  sulphu- 
retted hydrogen.  This  latter  then  reacts  upon  the  stannic  chloride, 
with  precipitation  of  stannic  sulphide — 

(1)  3SnCl2  +  6HC1  +  H2SO3  =•-  3SnCl4  +  3H2O  4-  H2S 

(2)  SnCl4  +  2H2S  =  SnS2  +  4HC1 

Nascent  •  hydrogen,*  obtained  by  the  action  of  hydrochloric  acid 
upon  zinc,  also  reduces  sulphurous  acid  to  sulphuretted  hydrogen, 

H2SO3  +  3H2  =  3H20  +  H2S 

The  test  may  be  made  by  adding  a  minute  trace  of  sulphurous  acid 
(or  a  solution  of  a  sulphite)  to  a  mixture  of  zinc  and  hydrochloric 
acid  in  a  test-tube,  and  applying  acetate  of  lead  paper  to  the  mouth 
of  the  tube. 

Sulphites.— The  only  sulphites  soluble  in  water  are  those  of 
the  alkali  metals.  They  are  all  decomposed  by  dilute  acids,  with 
evolution  of  sulphur  dioxide  (see  above).  Oxidising  agents  convert 
them  into  sulphates. 

When  heated  by  themselves,  most  sulphites  are  converted  into 
sulphides  and  sulphates— 

4K2SO3  =  K2S  +  3K2SO4 

Those  of  the  alkaline  earths  leave  an  oxide,  and  evolve  sulphur 
dioxide — 

BaSO3  =  BaO  +  SO2 

Barium  chloride  gives  a  white  precipitate  of  barium  sulphite, 
BaSO3,  soluble  in  dilute  HC1  (distinction  from  BaSO4). 

Lead  acetate  precipitates  white  lead  sulphite,  PbSO3.  The 
salt  undergoes  no  change  when  boiled  (contrast  Lead  thiosulphate). 

Silver  nitrate  gives  a  white  precipitate  of  silver  sulphite, 
which,  on  boiling,  is  converted  into  black  metallic  silver — 

Ag2SOs  +  H20  =  H2S04  +  Ag2 
A  sulphate  and  sulphite  in  solution  may  be  separated  by  acidulating 

*  Nascent  hydrogen,  obtained  by  the  action  of  sulphurous  acid  upon  zinc 
without  any  other  acid,  reduces  the  sulphurous  acid  to  hyposulphurous  (not  to 
be  confounded  with  thiosulphuric}  acid,  an  unstable  compound  which  passes 
into  thiosulphuric  acid  and  water  ;  thus — 

(1)  H2S03  +  H2  =  H20  +  H2S02 

(2)  2H2SO2  =  H2O  +  H2S2O3 


142  Qualitative  Analysis. 

the  dilute  solution  and  adding  barium  chloride.  The  precipitated 
sulphate  (insoluble  in  acid)  is  removed  by  nitration.  To  the  solu- 
tion, which  now  contains  barium  chloride  and  sulphurous  acid,  an 
oxidising  agent,  such  as  chlorine  water,  is  added,  when  a  precipitate 
of  barium  sulphate  is  again  thrown  down — 

BaCJ2  +  H2SO3  +  H2O  +  C12  =  BaSO4  +  4HC1 

Thiosulphuric  Acid  and  Thiosulphates. 

The  acid  is  unknown  in  the  free  state  ;  when  liberated  from  its 
salts  it  at  once  breaks  down  into  water,  sulphur  dioxide,  and  sulphur. 

Thiosulphates  of  barium,  lead,  and  silver  are  sufficiently  difficult 
of  solution  in  water,  to  be  precipitated  on  the  addition  of  a  solution 
of  a  thiosulphate  to  solutions  of  the  respective  metals.  All  others 
are  more  easily  soluble.  They  are  all  decomposed  by  acids,  with 
precipitation  of  sulphur  and  evolution  of  sulphur  dioxide  (this  is 
characteristic,  and  distinguishes  thiosulphates  from  sulphites,  which 
only  give  the  dioxide  without  depositing  sulphur}. 

When  heated  alone,  thiosulphates  are  converted  into  sulphates 
and  polysulphides. 

Barium  chloride  precipitates  from  a  moderately  strong  solu- 
tion white  barium  thiosulphate,  BaS2O3,  decomposed  by  hydro- 
chloric acid  ;  thus — 

BaS2O3  +  2HC1  =  BaCl2  +  H2O  +  SO2  +  S 

Silver  nitrate  gives  a  white  precipitate  of  silver  thiosulphate 
(easily  soluble  in  excess  of  sodium  thiosulphate,  forming  the  double 
salt,  AgNaS2O3),  which  quickly  changes  colour  owing  to  the  for- 
mation of  black  silver  sulphide  (contrast  the  behaviour  of  silver 
sulphite  above) — 

Ag2S203  +  H20  =  H2SO,  +  Ag2S 

Lead  acetate  behaves  in  a  similar  manner,  giving  black  lead 
sulphide  (contrast  lead  sulphite). 

Thiosulphates  are  oxidised  by  chlorine,  bromine,  and  other 
oxidising  agents,  in  the  same  manner  as  the  sulphites — 

NaaSaOa  +  4C12  +  5H2O  =  8HC1  +  2HNaSO4 
and  are  also  reduced  by  nascent  hydrogen,  with  formation  of  sul- 
phuretted hydrogen— 

Na2S203  +  2HC1  +  4H2  -  2NaCl  +  3H2O  +  2H2S 
Separation  of  the  Sulphur  Acids.— A  solution  containing 


Separation  of  the  Sulphur  Acids.  143 

a  sulphide,  sulphate,  sulphite,  and  thiosulphate  may  be  examined  in 
the  following  way  : — 

(1)  The   sulphide  is   first   separated  as  an   insoluble  metallic 
sulphide    in   such  a  way  as  to  avoid  the  liberation   of  any  acid, 
which  would   decompose   the  sulphite  or   thiosulphate.      This  is 
done  either  by  shaking  up  the  solution  with  a  little  lead  carbonate 
(or  cadmium  carbonate),  or  by  adding  a  solution  of  zinc  chloride 
which  has  been  rendered  alkaline  with  ammonia.     The  precipitated 
sulphide  is  then  removed  by  filtration.     Very  small  traces  of  sul- 
phuretted hydrogen  will  produce  a  distinct  coloration  in  the  white 
carbonate  of  lead. 

(2)  Barium  chloride  is  added  to  the  filtrate,  which  precipitates 
barium  sulphate  and  sulphite. 

(3)  The  filtrate  from  this  precipitate  is  acidified  with  hydro- 
chloric acid  and  warmed.    The  thiosulphate  is  thereby  decomposed, 
with  precipitation  of  sulphur  and  evolution  of  sulphur  dioxide. 

(4)  The  mixed  barium  sulphate  and  barium  sulphite  is  treated 
with   hydrochloric   acid.     The   sulphate  is  left,  and  the  sulphite 
dissolves.     On  filtering  and  adding  chlorine  water,  the  sulphurous 
acid  is  oxidised  to  sulphuric  acid,  and  a  precipitate  of  barium  sul- 
phate is  produced. 


CHAPTER  XIII. 

NITROGEN  AND  PHOSPHORUS. 

THE  properties  of  the  element  nitrogen  in  its  free  state  are  not 
employed  in  the  analytical  examination  of  its  compounds.* 

Nitric  Acid  and  Nitrates. 

Nitric  acid  is  a  fuming  corrosive  liquid,  miscible  with  water. 
It  readily  dissolves  most  metals,  converting  them  into  nitrates  or 
oxides,  with  evolution  of  oxides  of  nitrogen,  and  in  some  cases  with 
the  formation  of  ammonia  ;  e.g. — 

[Sn  +  4HNO3  =  H2SnO3  +  H2O  +  4NO2]  x  5  (see  Meta- 

stannic  acid) 

3Hg  +  8HN03  =  3Hg(N03)2  +  4H2O  +  2NO 
4Zn  +  ioHNO3  =  4Zn(NO3)2  +  5H2O  +  N2O 

4Zn  +  9HNO3  =  4Zn(NO3)2  +  sH2O  +  NH3 

The  reduction  of  the  nitric  acid  may  be  regarded  as  due  to  the 
action  of  nascent  hydrogen,  which  is  developed  by  the  first  action 
of  the  acid  upon  the  metal — 

2HNO3  +  3H2  =  4H2O  +  2NO 

The  extent  to  which  the  reduction  of  the  acid  will  take  place 
depends  upon  a  number  of  conditions — namely,  the  particular  metal, 
the  strength  of  acid,  the  temperature,  and  the  amount  of  nitrate 
which  is  present  in  the  mixture. 

*  The  presence  of  nitrogen  in  organic  compounds  may  readily  be  detected 
by  heating  a  small  quantity  of  the  solid  substance  in  a  test-tube  with  a  fragment 
of  the  metal  sodium  (or  potassium).  Under  these  circumstances  the  nitrogen 
unites  with  the  carbon  of  the  organic  compound,  forming  cyanogen,  which,  in 
the  presence  of  the  alkali  metal,  gives  a  cyanide  of  the  metal.  The  still  hot 
tube  should  be  plunged  into  a  small  quantity  of  water  in  a  beaker,  which 
immediately  breaks  it  up  and  allows  the  cyanide  to  dissolve,  and  the  excess  of 
sodium  to  be  quietly  acted  upon  by  the  water.  Two  or  three  drops  of  both 
ferrous  and  ferric  salts  are  added,  and  the  liquid  then  acidified  with  hydrochloric 
acid,  when  Prussian  blue  will  be  formed. 


Nitric  Acid  and  Nitrates.  145 

With  magnesium,  and  under  special  conditions  with  zinc  also, 
free  hydrogen  is  evolved  by  the  action  of  nitric  acid. 

Nitric  acid  also  oxidises  many  of  the  non-metals  ;  thus  sulphur, 
phosphorus,  and  iodine,  are  converted  respectively  into  sulphuric, 
phosphoric,  and  iodic  acids. 

Nitrates  are  all  soluble  in  water  ;  their  recognition,  therefore, 
is  based  upon  the  oxidising  reactions  of  which  they,  or  the  nitric 
acid  which  they  yield,  are  capable. 

Reduction  by  Ferrous  Salts. — When  ferrous  sulphate  is 
brought  into  contact  with  a  mixture  of  a  nitrate  and  strong  sulphuric 
acid,  the  solution  assumes  a  deep  brown  colour.  Three  chemical 
changes  go  to  make  up  the  reaction  :  (i)  the  liberation  of  nitric 
acid  by  the  action  of  sulphuric  acid  upon  the  nitrate  ;  (2)  the 
oxidation  of  the  ferrous  to  a  ferric  salt,  with  elimination  of  nitric 
oxide  ;  and  (3)  the  absorption  of  the  nitric  oxide  so  formed  by  a 
further  portion  of  ferrous  salt,  forming  an  unstable  brown  compound 
having  the  composition  NO,2FeSO4 — 

(1)  KN03  +  H2S04  -  HKS04  +  HNO3 

(2)  2HNO3  +  6FeSO4  +  3H2SO4  =  3Fe2(SO4)3  +  4H2O  +  2NO 
The  test  is  extremely  delicate,  and  is  carried  out  in  the  following 
manner.     The  solution  of  the  nitrate  is  mixed  with  about  its  own 
volume  of  strong   sulphuric  acid  in  a  test-tube,  and  the  mixture 
cooled.    To  this  a  little  ferrous  sulphate  solution  is  cautiously  added, 
the  tube  being  held  in  an  inclined  position,  so  that  the  ferrous  sul- 
phate shall  float  upon  the  denser  liquid  already  in  the  tube.    Where 
the  two  liquids  meet,  the  brown  colour  will  be  developed.     By  a 
gentle  movement  of  the  tube,  so  as  to  cause  a  slight  admixture  of 
the  liquids  at  the  point  where  they  meet,  the  brown  ring  will  be  still 
more  apparent.     The  coloured  compound  is  decomposed  by  heat, 
with  evolution  of  nitric  oxide,  hence  the  necessity  for  making  the 
test  with  cold  solutions.* 

Reduction  by  Sulphurous  Acid.— When  copper  (or  mer- 
cury) is  heated  with  sulphuric  acid  in  the  presence  of  a  nitrate,  nitric 
oxide  is  evolved,  which,  in  contact  with  the  air,  gives  red  vapours 
of  nitrogen  peroxide — 

3Cu  +  4H2S04  +  2KN03  =  3CuSO4  +  K2$O4  +  4H2O  +  2NO 
The  sulphur  dioxide  (developed  by  the  action  of  the  acid  upon  the 

*  Ferrous  chloride  in  the  presence  of  hydrochloric  acid  gives  a  precisely 
similar  result  with  a  nitrate ;  nitric  oxide  is  liberated,  which  is  absorbed  by  the 
ferrous  chloride  to  a  brown  liquid— 

KNO3  +  3FeCl2  +  4HC1  -=  3FeCl3  +  KC1  +  2H?O  +  NO 

L 


146  Qualitative  Analysis. 

copper)  is  oxidised  by  the  nitric  acid  (simultaneously  generated  by 
the  action  of  the  acid  upon  the  nitrate)  to  sulphuric  acid  ;  thus  — 


(1)  Cu  +  2H2SO4  =  CuSO4  +  2H2O  + 

(2)  2HNO3  +  sSO2  +  2H2O  ='  3H2SO4  +  2NO 


SO 


The  nitrate  is  mixed  with  a  little  strong  sulphuric  acid,  and  a  few 
fragments  of  copper  foil  or  turnings  are  introduced.  On  boiling 
the  mixture,  red  fumes  of  nitrogen  peroxide,  NO2,  will  appear  in  the 
tube,  which  will  be  more  easily  seen  by  looking  down  through  the 
mouth  of  the  tube.* 

4  A  similar  reduction  takes  place,  with  the  liberation  of  nitric  oxide, 
when  sulphur  dioxide  is  generated  in  presence  of  a  nitrate,  from 
either  a  sulphite  or  a  thiosulphate,  by  the  action  of  sulphuric  acid. 

A  small  crystal  of  sodium  thiosulphate  is  added  to  the  solution 
of  a  nitrate,  and  a  few  drops  of  strong  sulphuric  acid  added.  On 
gently  warming  the  mixture,  brown  fumes  are  seen  in  the  tube. 

Reduction  by  Nascent  Hydrogen,  (i)  With  Formation  of 
Nitrite.  —  When  a  nitrate  in  solution  is  exposed  to  the  gentle  action 
of  nascent  hydrogen—  derived  by  the  action  of  sodium  amalgam, 
zinc  amalgam,  or  copper-zinc  couple  —  the  nitrate  is  reduced  to 
nitrite  — 

KNO3  +  H2  =  H2O  +  KNO2 

This  test  is  very  delicate,  and  may  be  carried  out  as  follows  : 
A  small  piece  of  zinc  foil  (or  granulated  zinc)  is  placed  in  the  solution 
of  the  nitrate  in  a  test-tube,  and  one  drop  of  copper  sulphate  added 
(this  causes  the  deposition  of  a  minute  quantity  of  copper  upon  the 
zinc,  thus  creating  the  "  copper-zinc  couple  ").  The  mixture  is 
gently  boiled  for  a  minute  or  two.  One  drop  of  the  liquid  (after 
cooling)  is  placed  upon  a  piece  of  potassium-iodide-and-starch 
paper,f  and  then  touched  with  a  glass  rod  moistened  with  dilute 
sulphuric  acid.  The  paper  will  be  instantly  stained  blue  by  the 

*  The  mechanism  of  this  reaction  is  sometimes  explained  by  supposing  that 
the  nitric  acid,  set  free  from  the  nitrate,  acts  upon  the  copper  according  to  the 
familiar  equation  sCu  -f-  8HNO3  =  3Cu(NO3)2  +  4H2O  +  2NO.  But,  as  a 
matter  of  fact,  copper  sulphate,  not  nitrate,  is  found  in  solution,  the  whole  of 
the  nitrogen  being  converted  into  nitric  oxide.  Moreover,  charcoal  may  be 
substituted  for  the  copper.  If  dilute  nitric  acid  be  boiled  with  charcoal  no 
brown  fumes  are  formed,  but  on  the  addition  of  a  little  sulphuric  acid  they  at 
once  appear,  owing  to  the  action  of  the  sulphur  dioxide  which  is  evolved  from 
carbon  and  sulphuric  acid. 

f  Potassium-iodide-and-starch  paper  is  made  by  dipping  ordinary  white 
note  paper  into  moderately  thin  starch  containing  a  little  potassium  iodide,  and 
allowing  it  to  dry.  It  may  be  preserved  indefinitely  if  kept  in  a  stoppered 
bottle,  and  thus  obviates  the  necessity  of  constantly  making  a  little  starch  paste. 
Every  student  should  prepare  for  himself  a  stock  of  the  paper. 


Nitrous  Acid  and  Nitrites.  147 

liberation  of  iodine  and  formation  of  iodide  of  starch  (test  for  a 
nitrite). 

(2)  With  Formation  of  Ammonia, — By  the  prolonged  action  of 
the  copper-zinc  couple  ;  or  by  means  of  nascent  hydrogen  produced 
by  boiling  a  solution  of  caustic  soda  with  zinc,  nitrates  (and 
nitrites)  are  reduced  to  ammonia — 

KNO3  +  4H2  =  KHO  +  2H2O  +  NH3 

A  few  drops  of  the  solution  of  a  nitrate  are  added  to  caustic  soda 
in  a  test  tube,  along  with  a  little  granulated  zinc.  On  boiling, 
ammonia  is  evolved,  which  may  be  detected  in  the  usual  way. 

Indigo. — Nitric  acid  oxidises  indigo  blue(C8H5NO),  converting 
it  into  yellow  isatin  (C8H5NO2).  A  little  concentrated  sulphuric 
acid  is  tinted  with  a  minute  quantity  of  indigo,  and  a  few  drops  of 
a  solution  of  a  nitrate  added.  On  heating  the  mixture  the  blue 
colour  disappears.  (The  yellow  of  the  isatin  is  scarcely  visible  in 
small  quantities,  hence  the  solution  appears  bleached.) 

Nitrates  all  undergo  decomposition  when  strongly  heated. 
Nitrates  of  alkali  metals  and  alkaline  earths,  when  gently  heated, 
are  reduced  to  nitrites,  with  evolution  of  oxygen. 

Ammonium  nitrate  passes  into  water  and  nitrous  oxide.  Other 
nitrates,  e.g.  lead  nitrate,  leave  an  oxide  of  the  metal,  and  give  oft" 
oxygen  and  nitrogen  peroxide. 

When  heated  with  oxidisable  substances  (carbon,  sulphur,  etc.) 
the  decomposition  is  propagated  with  explosive  violence.  Thus, 
when  nitrates  are  heated  before  the  blowpipe  on  charcoal,  deflagra- 
tion of  the  charcoal  takes  place. 

Nitrates  and  chlorates,  when  present  ^together,  are  examined  by 
being  first  converted  by  heat  into  nitrites  and  chlorides.  If  present 
as  salts  of  metals  other  than  the  alkalies,  sodium  carbonate  is 
added,  and  the  dry  mixture  heated  until  the  evolution  of  oxygen  is 
at  an  end.  The  residue  is  extracted  with  water,  and  the  solution 
examined  for  nitrites  and  chlorides. 

If  chlorides  are  originally  present  as  well  as  nitrates  and 
chlorates,  they  must  be  first  removed  by  precipitation  with  silver 
sulphate,  as  explained  on  p.  133. 

Nitrous  Acid  and  Nitrites. 

The  acid  is  not  known  in  the  pure  state.  Even  when  liberated 
in  dilute  solutions,  it  speedily  breaks  up  into  nitric  acid,  nitric  oxide 
and  water.  Hence  when  nitrites  are  decomposed  by  acids,  nitric 


148  Qualitative  Analysis. 

oxide  is  evolved,  which,  in  contact  with  atmospheric  oxygen,  passes 
into  the  brown  gas  NO2 ;  thus — 

6NaNO2  +  3H2SO4  =  3Na2SO4  +  2HNO3  +  2H2O  +  4NO 

Nitrites  are  all  soluble  in  water,  but  the  silver  salt  is  suffi- 
ciently difficult  of  solution  to  be  precipitated,  on  the  addition  of 
silver  nitrate,  to  a  (not  too  dilute)  solution  of  a  nitrite. 

All  nitrites  are  easily  decomposed  by  dilute  acids  in  the  cold, 
with  evolution  of  nitric  oxide,  as  shown  in  the  above  equation.  If 
the  action  takes  place  in  the  presence  of  a  ferrous  salt,  the  same 
brown-coloured  compound  is  produced  as  in  the  case  of  a  nitrate. 
[Nitrites  therefore  give  a  "brown  ring,"  when  dilute  sulphuric,  or 
even  acetic,  acid  is  used  (distinction  from  nitrates}.'] 

Oxidation  Reactions. — Nitrous  acid  and  nitrites  part  with 
oxygen,  and  are  converted  into  nitric  oxide — 

2HNO2  =  2NO  +  H2O  +  O 

Thus,  when  a  nitrite  is  acidified  with  dilute  sulphuric  acid  in  the 
presence  of  potassium  iodide,  the  nitrous  acid  first  formed  oxidises 
the  potassium  iodide  (or  hydriodic  acid),  setting  free  the  iodine  : 
the  liberated  iodine  is  detected  by  its  action  upon  starch — 

2K1  +  H20  +  O  =  2KHO  +  I2 

A  few  drops  of  a  solution  of  a  nitrite  are  placed  upon  potassium 
iodide-and-starch  paper,  and  a  single  drop  of  very  dilute  sulphuric 
acid  added  (by  means  of  a  glass  rod  dipped  in  the  acid)  when  a 
blue  stain  at  once  appears  upon  the  paper.  [Or  starch  emulsion 
may  be  added  to  the  solution  of  the  nitrite,  then  a  drop  or  two  of 
potassium  iodide,  and  lastjy  a  small  quantity  of  acid.] 

Sulphuretted  hydrogen  is  similarly  oxidised  by  a  nitrite  in 
presence  of  an  acid,  with  precipitation  of  sulphur — 

H2S  +  2HNO2  =  2H2O  +  2NO  +  S 

Nitrites  of  the  alkali  metals  are  decomposed  by  sulphuretted 
hydrogen  without  the  addition  of  an  acid,  giving  alkaline  sulphides. 

Reduction  Reactions.  —  Nitrous  acid,  by  absorption  of 
oxygen,  passes  into  nitric  acid  ;  it  therefore  is  capable  of  reducing 
other  compounds,  such  as  chromates,  permanganates,  mercurous 
(but  not  mercuric)  salts,  and  gold  compounds.  Potassium  per- 
manganate (in  presence  of  acid)  is  converted  into  manganous  salts, 
the  violet  colour  of  the  permanganate  being  destroyed.  Gold  and 
mercurous  salts  are  reduced  to  the  respective  metals  ;  thus— 


Phosphoric  Acid  and  Phosphates.  149 

Hg2G2  +  H20  +  HN02  =  HN03  +  2HC1  +  2Hg 
2AuCl3  +  3H2O  +  3HNO2  =  3HNO3  +  6HC1  +  2Au 

Detection  of  Nitrites  and  Nitrates  in  the  Same  Solu- 
tion.— Owing  to  the  ready  decomposition  of  nitrites  by  dilute 
acids,  they  are  easily  detected  in  presence  of  nitrates,  either  by  the 
liberation  of  iodine,  the  oxidation  of  ferrous  salts,  or  reduction  of 
potassium  permanganate.  To  find  a  nitrate  when  nitrites  are 
present  is  less  simple.  The  dilute  solution  of  the  mixed  nitrate 
and  nitrite  is  acidified  with  three  or  four  drops  of  dilute  sulphuric 
acid,  and  a  little  ferrous  sulphate  solution  (or  a  small  crystal  of  the 
salt)  is  added.  The  solution  at  once  becomes  dark  brown  (owing 
to  the  absorption,  by  the  ferrous  salt,  of  the  nitric  oxide  liberated 
from  the  nitrite).  It  is  then  heated  (but  not  allowed  to  boil),  with 
frequent  shaking,  when  nitric  oxide  is  expelled,  and  the  liquid 
gradually  becomes  colourless.  The  mixture  is  cooled,  and  one  drop 
more  dilute  acid  added,  and  a  little  more  ferrous  sulphate.  (If  all 
the  nitrite  present  has  been  decomposed,  this  addition  gives  no 
further  coloration.)  This  solution  is  now  poured  carefully  on  to  a 
small  quantity  of  strong  sulphuric  acid  in  a  test-tube,  so  as  to  float 
upon  the  acid,  and  where  the  liquids  meet  a  "  brown  ring  "  will  be 
formed,  due  to  the  nitrate  present. 

Phosphorus. 

The  property  which  phosphorus  possesses  of  emitting  a  feeble 
luminosity  when  exposed  to  the  air,  either  in  the  solid  state  or 
when  vaporised,  is  made  use  of  in  analysis  for  the  detection  of  the 
element  in  the  free  state.  The  substance,  mixed  with  water,  is 
boiled  in  a  flask  with  a  narrow  neck  in  a  dark  room.  The  issuing 
steam  will  then  appear  luminous  as  it  escapes  from  the  flask  ;  and 
presents  the  appearance  of  a  lambent  pale  greenish  flame.  A  piece 
of  paper  moistened  with  silver  nitrate,  when  held  over  the  neck  of 
the  flask,  will  become  blackened,  owing  to  the  formation  of  silver 
phosphide.* 

Phosphoric  Acid  and  Phosphates. 

Three  phosphoric  acids  (each  with  its  series  of  phosphates)  are 
known,  namely,  orthophosphoric  acid,  H3PO4 ;  pyrophosphoric 
acid,  H4P2O7  ;  and  metaphosphoric  acid,  HPO3. 

*  Phosphorus  is  a  violent  poison,  and  it  is  practically  only  in  cases  of 
toxicological  examination  that  free  phosphorus  is  sought  for.  The  allotropic 
form  ("red  phosphorus")  is  not  poisonous.  It  is  changed  by  heat  into  the 
ordinary  variety. 


150  Qualitative  Analysis. 

Orthophosphoric  acid  is  tribasic,  and  hence  produces  three 
classes  of  orthophosphates,  by  the  replacement  of  one,  two,  or 
three  of  the  hydrogen  atoms  ;  e.g. — 

Ammonium  magnesium  phosphate,  (NH4)MgPO4. 

Hydrogen  sodium  ammonium  phosphate  (microcosmic  salt], 
HNa(NH4)PO4. 

Dihydrogen  sodium  phosphate,  H2NaPO4. 

Normal  orthophosphates  (not  containing  a  volatile  base,  as 
NH4)  are  not  decomposed  when  heated  alone  ;  those  containing 
one  or  two  hydrogen  atoms,  or  ammonium  as  the  base,  are  con- 
verted into  pyro  or  meta  salts  (see  below).  The  only  orthophos- 
phates which  are  soluble  in  water  are  those  of  the  alkali  metals. 

Silver  nitrate  gives  a  yellow  precipitate  with  soluble 
phosphates,  of  silver  phosphate,  Ag3PO4,  which  distinguishes  ortho 
from  pyro  and  meta  compounds. 

The  chief  analytical  reactions  of  the  orthophosphates  have 
already  been  considered  in  Chap.  VII.,  p.  63. 

Pyro  phosphoric  Acid  and  Pyrophosphates.  —  When 
orthophosphoric  acid,  or  a  phosphate  containing  either  one  hydro- 
gen atom  or  one  ammonium  radical,  is  heated,  it  loses  water,  and 
is  converted  into  pyrophosphoric  acid  or  a  salt ;  thus — 

2H3PO4  =  H2O  +  H4P2O7 
2HNa2PO4  =  H2O  +  Na4P2O7 
2NH4MgP04  =  H20  +  2NH3~  +  Mg2P2O7 

Boiling  with  acids  retransforms  pyrophosphates  into  ortho- 
phosphates. 

Only  the  pyrophosphates  of  the  alkalies  are  soluble  in  water. 

Silver  nitrate  gives  a  white  precipitate  of  silver  pyrophos- 
phate,  Ag4P2O7. 

Magnesium  sulphate  precipitates  white  magnesium  pyro- 
phosphate,  Mg2P2O7,  soluble  in  excess  of  magnesium  sulphate,  and 
not  reprecipitated  in  the  cold  by  ammonia  (distinction  from  ortho- 
phosphates}. 

Ammonium  molybdate  gives  no  precipitate  until  the  "pyro" 
acid  has  been  changed  to  "  ortho  "  by  the  action  of  the  nitric  acid 
present. 

Metaphosphoric  Acid  and  Metaphosphates.f— The  acid 
is  formed  when  the  "  ortho  "  or  "  pyro  "  acids  are  strongly  heated, 
whereby  water  is  expelled.  It  is  known  as  glacial  phosphoric 
acid — 

H3P04  =  H20  +  HP03 

*  Metaphosphoric  acid  has  the  curious  property  of  forming  a  number  of 
salts  which  may  be  regarded  as  polymers  of  the  ordinary  salts  (see  Newth's 
"  Inorganic  Chemistry,"  p.  477). 


Hypophosphorous  Acid.  151 

The  reaction  is  reversible  ;  for  when  metaphosphoric  acid  is  dis- 
solved in  water,  it  passes  back  to  the  ortho  acid  —  slowly  in  the 
cold,  quickly  when  boiled. 

Silver  nitrate  gives  a  white  precipitate  of  silver  metaphos- 
phate,  AgPO3. 

Magnesium  sulphate,  in  presence  of  ammonium  chloride, 
gives  no  precipitate  (distinction  from  "  pyro  "  and  "  ortho  "  acids). 

Albumen  (white  of  egg)  is  coagulated  when  shaken  up  with 
metaphosphoric  acid  (or  metaphosphates  acidified  with  acetic  acid) 
(distinction  from  "  pyro  "  and  "  ortho  "  acids,  which  are  without  action 
upon  albumen]. 

Phosphorous  Acid  and  Phosphites. 

Phosphorous  acid,  H3PO3,  is  tribasic,  but  the  most  stable  salts 
are  those  which  contain  one  atom  of  hydrogen.  Thus,  trisodium 
phosphite,  Na3PO3,  is  decomposed  by  water  into  hydrogen  disodium 
phosphite,  HNa2PO3.  With  the  exception  of  phosphites  of  the 
alkalies,  the  salts  of  this  acid  are  either  insoluble  in  water  or  dis- 
solve only  with  difficulty,  and  may  therefore  be  precipitated  by 
double  decomposition  ;  thus,  on  the  addition  of  -barium  chloride 
to  a  solution  of  sodium  phosphite,  barium  phosphite,  HBaPO3,  is 
thrown  down  as  a  white  precipitate.  Both  the  acid  and  its  salts 
are  powerful  reducing  agents  ;  thus,  with  silver  nitrate  a  white 
precipitate  momentarily  forms,  which  quickly  becomes  black  from 
reduced  silver.  Similarly,  mercuric  compounds  are  converted  first 
to  mercurous  and  then  to  metallic  mercury.  When  strongly  heated, 
phosphites  decompose,  and  give  off  phosphoretted  hydrogen. 

Hypophosphorous  Acid  and  Hypophosphites. 

Hypophosphorous  acid,  H3PO2,  is  monobasic  ;  its  formula  may 
therefore  be  written  H(H2PO2).  The  sodium  salt  has  the  com- 
position Na(H2PO2),  while  the  barium  salt  is  expressed  by  the 
formula  Ba(H2PO2)2.  All  the  salts  are  soluble  in  water,  therefore 
barium  chloride  gives  no  precipitate  with  solutions  of  hypophos- 
phites  (distinction  from  phosphites). 

Hypophosphorous  acid  and  its  salts  are  still  more  powerful 
reducing  agents  than  phosphites.  One  characteristic  reaction  of 
this  nature,  which  distinguishes  hypophosphites  from  phosphites, 
is  the  reduction  of  copper  sulphate  to  copper  hydride,  Cu2H2. 

On  Adding  copper  sulphate  to  an  acidulated  solution  of  a 
hypophosphite,  and  gently  warming  the  mixture,  a  precipitate  is 
obtained  —  at  first  yellowish-brown,  quickly  turning  to  a  dark 
chocolate-brown  colour  ;  thus  — 


This  precipitate  of  cuprous  hydride  is  distinguishable  from  the  red 
cuprous  oxide,  not  only  by  its  much  darker  colour,  but  by  the 


152  Qualitative  Analysis. 

fact  that  when  treated  with  strong  hydrochloric  acid,  it  evolves 
hydrogen  ;  thus  — 

Cu2H2  +  2HC1  =  Cu2Cl2  +  2H2 

Under  the  influence  of  nascent  hydrogen  from  zinc  and  hydro- 
chloric acid,  hypophosphites  as  well  as  phosphites  give  phospho- 
retted  hydrogen,  which  imparts  to  the  hydrogen,  when  it  is  inflamed, 
the  characteristic  colour  of  the  phosphorus  flame. 

Hypophosphorous  acid  and  its  salts,  when  gently  heated,  give 
off  phosphoretted  hydrogen,  leaving  phosphoric  acid — 

2H3P02  =  H3P04  +  PH3 


CHAPTER    XIV. 
CARBON,  SILICON,  BORON. 

Carbon. 

THE  properties  of  carbon  by  which  it  is  most  readily  recognised 
are  (i)  its  combustibility  in  air  or  oxygen,  with  formation  of  only 
carbon  dioxide  ;  and  (2)  its  power  to  withstand  the  action  of  chlorine 
even  at  high  temperatures,  which  distinguishes  it  from  all  metals 
or  black  substances. 

Carbonic  Acid  and  Carbonates. 

Carbonic  Acid,  H2CO3,  is  an  unstable  compound  only  capable 
of  existence  in  dilute  aqueous  solution.  It  is  formed  when  carbon 
dioxide  is  dissolved  in  water,  and  has  a  feeble  acid  reaction.  It  is 
capable  of  dissolving  the  normal  carbonates  of  the  alkaline  earths 
and  of  magnesium,  forming  the  so-called  "acid"  carbonates  or 
"  bicarbonates "  (the  pharmaceutical  preparation  known  as  fluid 
magnesia  is  a  case  in  point) — 

MgC03  +  H2C03  =  H2Mg(C03)2 

All  normal  carbonates  are  insoluble  in  water,  except  those  of  the 
alkalies.  The  "  acid "  carbonates  are  all  soluble,  but  on  boiling 
their  solutions,  they  are  converted  into  normal  salts,  which  (except 
in  the  case  of  the  alkaline  carbonates)  are  then  precipitated  ;  thus  — 

H2Ca(C03)2  =  CaC03  +  H2O  +  CO2 

(The  formation  of  boiler  incrustations  and  the  "  furring  "  of  kettles 
are  due  to  this  decomposition.) 

The  normal  carbonates  of  the  alkalies  are  by  this  reaction 
readily  distinguished  from  the  bicarbonates.  Thus,  on  heating  a 
solution  of  sodium  bicarbonate,  effervescence  rapidly  sets  in  owing 
to  the  escape  of  carbon  dioxide — 

2HNaCO3  =  Na2CO3  +  H2O  4-  CO2 
These  two   classes   of  salts   may   also    be   distinguished    by   the 


154 


Qualitative  Analysis. 


difference  in  their  behaviour  towards  certain  metallic   solutions } 
e.g.  magnesium  or  mercuric  salts.     Thus — 

With  magnesium  sulphate  or  chloride,  normal  sodium 
carbonate  gives  a  white  precipitate  of  a  basic  carbonate — 

5MgSO4  +  5Na2CO3  =  MgO,4MgCO3  +  sNa2SO4  +  CO, 
Sodium  bicarbonate  gives  no  precipitate  with  magnesium  salts,  the 
reason  being  that  in  the  presence  of  the  large  quantity  of  carbonic 
acid  set  at  liberty  from  the  bicarbonate,  the  magnesium  carbonate 
is  redissolved,  forming  the  soluble  magnesium  bicarbonate. 

With  mercuric  chloride,  the  normal  carbonate  gives  a 
reddish  precipitate  of  a  basic  oxide,  while  the  bicarbonate  gives 
no  precipitate. 

All  carbonates  are  decomposed  by  dilute  hydrochloric  acid 
(and  by  nearly  all  acids)  with  effervescence,  due  to  the  rapid  escape 
of  carbon  dioxide.  The  gas  is  identified  by  its  action  upon  lime- 
water  (or  baryta-water). 

The  test  is  made  by  adding  a  few  drops  of  acid  to  the  carbonate 
in  a  test-tube,  and  decanting  the  evolved  (heavy)  gas  into  a  second 
test-tube  containing  a  little  lime-water,  Ca(HO)2.  On  shaking  the 
lime-water  with  the  gas,  the  liquid  becomes  milky,  owing  to  the 
precipitation  of  calcium  carbonate. 

The  only  other  gas  which  gives  a  white  precipitate  with  lime- 
water  is  sulphur  dioxide.  This  is 
easily  distinguished  from  carbon 
dioxide  by  its  smell.  If  the  two 
gases  are  present  together,  the 
sulphur  dioxide  (recognised  by 
its  odour)  may  be  removed  by 
means  of  potassium  permanga- 
nate. The  two  gases  (liberated 
simultaneously  by  the  action  of 
an  acid  upon  a  mixture  of  a  car- 
bonate and  sulphite)  are  passed 
through  a  little  potassium  per- 
manganate solution  in  a  test-tube, 
fitted  as  shown  in  Fig.  14.  The 
sulphur  dioxide  is  absorbed  (being 
oxidised  by  the  permanganate  into 
sulphuric  acid),  and  the  carbon 
dioxide  passes  on,  and  can  be  de- 
tected by  means  of  lime-water. 
If  the  passage  of  the  gases  through 
FlG-  J4-  the  permanganate  be  continued 

for  a  few  minutes,  the  colour  of  the  solution  becomes  entirely 
destroyed  ;  and  the  liquid  may  then  be  tested  for  sulphuric  acid  in 
the  usual  way. 


Oxalic  Acid  and  Oxalates.  155 

When  strongly  heated,  the  normal  carbonates  of  the  alkali 
metals  (not  ammonium)  remain  unchanged.  Those  of  the  alkaline 
earths  are  converted  at  a  high  temperature  into  oxides,  with 
evolution  of  carbon  dioxide  (illustrated  in  the  process  of  lime- 
burning).  All  other  carbonates  are  more  readily  decomposed  by 
heat. 

Formic  Acid  and  Formates. 

Formic  acid,  H2CO2  or  H(COHO)  (the  acid  present  in  ants 
and  in  nettles),  is  a  pungent-smelling  colourless  liquid.  The  acid 
is  monobasic.  All  formates  are  soluble  in  water.  Both  the  acid 
and  the  salts  are  powerful  reducing  agents,  as  one  molecule  of  the 
acid  is  capable  of  withdrawing  one  atom  of  oxygen  (or  its  equivalent 
in  chlorine)  from  compounds  capable  of  reduction.  Thus,  mercuric 
chloride  is  reduced  to  mercurous  chloride,  and  carbon  dioxide  is 
eliminated-r- 

2HgCl2  +  H2C02  =  C02  +  2HC1  +  Hg2Cl2 

Formic  acid  and  formates  are  immediately  decomposed  by  concen- 
trated sulphuric  acid  without  the  application  of  heat,  into  water 
(absorbed  by  the  acid)  and  carbon  monoxide — • 

H(COHO)  =.H2O  +  CO 

The  escaping  carbon  monoxide  is  distinguished  by  burning  with  a 
fine  pale-blue  flame. 

Oxalic  Acid  and  Oxalates. 

Oxalic  acid,  H2C2O4,  is  a  white  solid,  soluble  in  water,  and 
crystallising  from  the  solution  with  two  molecules  of  water  of 
crystallisation,  H2C2O4,2H2O.  The  acid  is  dibasic. 

The  oxalates  of  the  alkalies  are  soluble  in  water,  as  well  as 
certain  acid  oxalates  (e.g.  acid  oxalate  of  barium).  All  other 
oxalates  are  insoluble  or  only  sparingly  soluble. 

Insoluble  oxalates  are  decomposed  by  mineral  acids,  with  the 
formation  of  salts  of  the  mineral  acid  with  the  metal  contained  in 
the  oxalate,  and  the  liberation  of  oxalic  acid.  In  the  case  of  strong 
sulphuric  acid,  the  liberated  oxalic  acid  is  itself  decomposed  on 
the  application  of  a  gentle  heat  ;  carbon  dioxide  and  monoxide 
being  disengaged  in  equal  volumes  ;  thus — 

H2C204  =  H20  +  C02  +  CO 

the  sulphuric  acid  in  this  instance,  as  in  the  case  of  formic  acid, 
taking  up  the  elements  of  water  from  the  molecule. 

In  the  presence  of  reducible  compounds,  oxalic  acid  passes  into 
water  and  carbon  dioxide  only,  one  molecule  of  the  acid  taking  up 


156  Qualitative  Analysis. 

one  atom  of  oxygen.  In  this  way  it  is  able  to  reduce  potassium 
permanganate,  two  molecules  of  which  furnish  five  atoms  of 
oxygen,  and  therefore  oxidise  five  molecules  of  oxalic  acid  ;  or  one 
molecule  of  permanganate  causes  the  liberation  of  five  molecules 
of  carbon  dioxide — 

2KMnO4  +  3H2SO4  =  K2SO4  +  2MnSO4  +  3H2O  +  50 
50  +  5H2C2O4  =  5H20  4-  ioCO2 

Similarly,  manganese  dioxide,  having  one  atom  of  available  oxygen, 
when  passing  to  the  condition  of  a  manganous  salt,  is  able  to 
oxidise  one  molecule  of  oxalic  acid  ;  thus — 

MnO2  +  H2SO4  +  H2C2O4  =  MnSO4  +  2H2O  +  2CO2 

Of  the  insoluble  oxalates,  the  calcium  salt,  CaC2O4,  is  of  the  most 
importance  analytically.  Its  formation  as  a  special  test  for 
calcium  has  already  been  considered  (p.  33).  The  precipitate  is 
only  slightly  soluble  in  oxalic  acid  (whereas  barium  oxalate  is 
somewhat  readily  dissolved  by  oxalic  acid,  forming  a  soluble  acid 
oxalate),  and  is  scarcely  dissolved  by  acetic  acid.  It  is,  however, 
readily  soluble  in  hydrochloric  acid. 

When  heated,  oxalates  are  all  decomposed.  Those  of  the 
alkalies  and  alkaline  earths  are  converted  into  carbonates,  with 
evolution  of  carbon  monoxide.  The  oxalates  of  metals  which 
either  do  not  form  carbonates,  or  whose  carbonates  are  decomposed 
by  heat,  leave  a  metallic  oxide,  and  give  off  carbon  dioxide  and 
monoxide  either  together  or  in  two  stages  ;  thus — 

(1)  CaC2O4  =  CaC03  +  CO 

(2)  CaC03  =  CaO  +  CO2 

In  the  case  of  metals  whose  oxides  are  decomposed  by  heat,  e.g. 
silver  oxalate,  the  metal  is  left. 

Acetic  Acid  and  Acetates. 

Acetic  acid,  C2H4O2  or  H(C2H3O2),  is  a  white  crystalline 
solid,  which  melts  at  the  temperature  of  a  warm  room  (16°  C). 
The  common  reagent  is  a  solution  of  the  acid  in  water.  The  acid 
is  monobasic.  The  acetates  are  all  soluble  in  water,  the  silver  and 
mercurous  salts  being  sufficiently  difficult  of  solution  to  allow  of 
their  being  precipitated  by  the  addition  of  moderately  strong  acetic 
acid  to  solutions  of  the  respective  metals.  Ferric  chloride,  when 
added  to  a  solution  of  an  acetate,  gives  ferric  acetate,  which  is  a 
soluble  salt  imparting  a  dark  red  colour  to  the  solution.  When  the 
liquid  is  boiled,  a  brownish  precipitate  separates  out,  consisting  of 


Tar  tar  ic  and  Citric  Acids.  157 

basic  ferric  acetate  (p.  69).  The  acetates  are  decomposed  by  mineral 
acids,  acetic  acid  being  liberated  ;  and  as  acetic  acid  is  volatile 
without  undergoing  decomposition,  it  can  be  distilled  off  and 
recognised  by  its  odour.  When  an  acetate  is  mixed  with  strong 
sulphuric  acid  and  a  little  alcohol,  (C2H5)HO,  and  the  mixture 
gently  heated,  a  reaction  takes  place,  resulting  in  the  formation  of 
ethyl  acetate  (acetic  ether],  a  volatile,  ethereal,  and  pleasant-smelling 
compound.*  The  mechanism  of  the  reaction  is  as  follows  :  (i) 
The  sulphuric  acid  acts  upon  the  alcohol,  forming  a  compound 
known  as  sulphethylic  acid,  H(C2H5)SO4 ;  and  (2)  this  then  reacts 
with  the  acetic  acid  liberated  by  the  sulphuric  acid  from  the 
acetate  ;  thus — 

(1)  H2S04  +  (C2H5)HO  -  H(C2HJS04  +  H2O 

(2)  H(C2H5)S04  +  H(C2H302)  -  H2S04  +  C2H5(C2H3O2) 

When  acetates  are  heated  alone,  they  blacken  slightly,  and  evolve 
vapours  of  a  volatile  liquid  known  as  acetone.  The  vapour  has  a 
characteristic  smell,*  and  is  inflammable. 

Tartaric  Acid  and  Tartrates. 

Tartaric  acid,  H2(C4H4O6),  is  a  white  crystalline  solid  ;  decom- 
posed by  heat,  and  therefore  not  volatile.  It  is  soluble  in  water. 
The  acid  is  dibasic. 

The  normal  tartrates  of  the  alkalies  are  freely  soluble  in  water ; 
the  "  acid  "  tartrates  (especially  of  potassium  and  ammonium  (see 
pp.  21  and  23)  are  sparingly  soluble.  Other  normal  tartrates  are  either 
insoluble  or  difficultly  soluble,  but  most  of  them  dissolve  in  tartaric 
acid,  forming  "  acid  "  salts.  The  tartrates  show  a  great  tendency 
to  form  soluble  double  salts  with  alkalies  ;  thus  the  tartrates  of  the 
metals  of  Group  III.  are  dissolved  by  alkalies,  owing  to  the  forma- 
tion of  double  tartrates.  Hence,  in  the  presence  of  tartaric  acid, 
alkalies  fail  to  give  precipitations  with  solutions  of  those  metals. 

Tartrates  are  decomposed  by  mineral  acids,  with  elimination  of 
tartaric  acid.  If  heated  with  strong  sulphuric  acid,  the  liberated 
tartaric  acid  is  broken  up  into  water,  carbon  dioxide,  carbon 
monoxide,  and  separation  of  carbon  (the  mixture  blackens  there- 
fore} ;  and  from  the  action  of  the  sulphuric  acid  upon  the  carbon, 
sulphur  dioxide  is  formed — 

H.2(C4H406)  =  3H20  +  C02  +  CO  +  2C 
C  +  2H2SO4  =  2H2O  +  CO2  +  2SO2 

Calcium  hydroxide  (lime-water),  added  to  a  solution  of 
tartaric  acid,  gives  no  precipitate,  unless  added  in  sufficient  excess 

*  No  description  of  odours  of  this  kind  can  convey  an  exact  idea  of  them. 
It  is  therefore  only  by  preparing  and  smelling  the  compound,  and  so  actually 
learning  its  characteristic  smell,  that  the  substance  can  afterwards  be  recog- 
nised by  such  a  test  as  this. 


158  Qualitative  Analysis. 

to  ensure  the  formation  of  the  normal  tartrate,  which  then  pre- 
cipitates out. 

Soluble  calcium  salts  (not  the  sulphate)  give  the  same  precipi- 
tate when  added  to  solutions  of  normal  tartrates.  The  precipitate 
is  soluble  in  tartaric  acid,  and  also  in  acetic  acid  (contrast  Calcium 
oxalate). 

Silver  nitrate  gives,  with  solutions  of  normal  tartrates,  a 
white  precipitate  of  silver  tartrate.  The  precipitate  is  soluble  in 
ammonia,  and  on  gently  warming  the  ammoniacal  solution,  the 
silver  is  precipitated  as  a  coherent  film  or  mirror  upon  the  glass 
vessel.  The  test  is  made  in  the  following  way  : — 

A  small  quantity  of  a  normal  tartrate  (the  double  tartrate  of 
sodium  and  potassium,  Rochelle  salt,  is  a  suitable  salt  to  employ 
in  order  to  study  the  reaction),  not  much  larger  than  a  pin's  head, 
is  dissolved  in  a  little  water  in  a  carefully  cleaned  test-tube,  and  a 
few  drops  of  silver  nitrate  added.  Dilute  ammonia  is  then  added 
drop  by  drop,  until  the  precipitated  silver  tartrate  is  nearly  wholly 
dissolved.  The  mixture  is  then  diluted  with  water,  so  as  to  about 
half  fill  the  test-tube.  The  tube  is  then  placed  into  boiling  water 
for  a  few  minutes,  when  the  silver  will  be  deposited  as  a  brilliant 
mirror  upon  the  glass.  ( This  reaction  distinguishes  tartaric  acid 
from  all  the  other  acids  treated  in  this  chapter?) 

When  heated  alone,  tartaric  acid  and  tartrates  are  decomposed, 
evolving  vapours  which  are  inflammable,  and  emitting  a  smell 
resembling  that  of  burnt  sugar.  The  residue,  in  the  case  of 
tartrates  of  the  alkalies  and  alkaline  earths,  consists  of  carbon  and 
the  carbonate  of  the  alkali  ;  while  other  tartrates  leave  either  a 
metallic  oxide  or  metal  mixed  with  carbon. 

Citric  Acid  and  Citrates. 

Citric  acid,  H3(C6H5O7),  closely  resembles  tartaric  acid  in 
appearance.  It  is  soluble  in  water.  The  acid  is  tribasic. 

The  citrates  of  the  alkalies  are  readily  soluble ;  those  of  the 
alkaline  earths  are  less  easily  soluble  ;  calcium  citrate,  however, 
being  soluble  in  cold  water,  but  nearly  insoluble  in  hot  water. 

Calcium  hydroxide,  or  calcium  chloride,  if  moderately 
dilute,  gives  no  precipitate  in  the  cold,  when  added  to  solutions 
of  citric  acid  or  normal  citrates.  (A  strong  solution  of  calcium 
chloride  mixed  with  a  strong  solution  of  a  citrate  may  give  a 
precipitate  even  in  the  cold.  With  more  dilute  solutions,  the 
calcium  citrate  is  gradually  precipitated  ;  completely  only  after  a 
few  hours)  ;  but  on  heating  the  mixture,  a  white  precipitate  of 
calcium  citrate  is  produced  (distinction  from  tartaric  add}.  The 
precipitate  is  soluble  in  acetic  acid  (distinction  from  calcium 
oxalate). 

By  means  of  the  reaction  with  calcium  chloride,  tartaric  and 
citric  acids  may  be  detected  when  mixed  together.  Calcium 
chloride  is  added  to  the  cold  and  not  too  concentrated  solution, 
whereby  calcium  tartrate  is  precipitated  ;  on  filtering  and  heating 
the  filtrate,  a  further  precipitation  takes  place  of  calcium  citrate. 


Hydrocyanic  Acid  and  Cyanides.  159 

Silver  nitrate  gives  a  white  precipitate  of  silver  citrate, 
soluble  (like  the  tartrate)  in  ammonia  ;  but  the  ammoniac  al  solu- 
tion does  not  deposit  metallic  silver  when  warmed  (compare  Silver 
tartrate). 

Cyanogen. 

Cyanogen,  (CN)  or  Cy,  is  a  negative  compound  radical,  which 
resembles  the  halogen  elements  in  many  of  its  chemical  habits. 
It  is  liberated  when  the  cyanides  of  many  of  the  heavy  metals  (e.g. 
mercuric  cyanide)  are  heated  alone. 

Cyanogen  is  a  colourless  gas,  with  an  odour  which  is  like  the 
taste  of  bitter  almonds  or  the  kernels  of  cherry-stones.*  It 
burns  with  a  characteristic  flame,  having  a  beautiful  mauve-pink 
colour.  The  gas  is  soluble  in  water,  but  the  solution  is  unstable. 
It  also  dissolves  in  yellow  ammonium  sulphide,  forming  ammonium 
thiocyanate — 

(NH4)2S2  +  2Cy  =  2NH4CyS 

This  reaction  forms  a  ready  method  for  the  detection  of  small 
quantities  of  cyanogen.  The  test  is  made  in  the  following  way. 
A  piece  of  filter-paper,  moistened  with  a  drop  of  yellow  ammonium 
sulphide,  is  held  over  the  mouth  of  the  test-tube  in  which  the 
compound  is  being  heated.  If  the  yellow  colour  of  the  sulphide  is 
destroyed,  the  colourless  spot  is  touched  with  a  glass  rod  dipped 
in  ferric  chloride,  when  a  red  stain  of  ferric  thiocyanate  is  at  once 
obtained.  If  the  quantity  of  cyanogen  is  so  small  that  the  colour 
of  the  ammonium  sulphide  is  not  destroyed,  a  single  drop  of  a 
solution  of  zinc  chloride  or  sulphate  is  placed  upon  it,  using  a 
pipette  (this  removes  the  excess  of  ammonium  sulphide,  forming 
white  zinc  sulphide),  after  which  the  spot  is  touched  with  ferric 
chloride,  when  the  red  stain  will  appear.  (Ferric  chloride  cannot 
be  added  until  the  ammonium  sulphide  is  removed,  or  a  black 
stain  of  ferrous  sulphide  will  be  produced.) 

Hydrocyanic  Acid  and  Cyanides. 

Hydrocyanic  acid,  HCN  or  HCy,  is  the  analogue  of  the 
hydrogen  acids  of  the  halogens.  In  the  pure  state  it  is  a  very 
volatile  liquid  (B.P.,  27°  C.).  It  is  liberated  when  certain  cyanides 
(see  below)  are  decomposed  by  acids,  and  may  be  recognised  by 

*  Cyanogen,  and  also  hydrocyanic  acid,  are  most  deadly  poisons,  and 
therefore  the  greatest  caution  must  be  exercised  in  smelling  either  of  these 
compounds.  The  odour  of  the  greatly  diluted  gas  seems  to  be  perceived  more 
by  the  sense  of  taste  than  of  smell,  being  recognised  as  a  somewhat  bitter  but 
not  unpleasant  sensation  at  the  back  of  the  throat. 


160  Qualitative  Analysis. 

its  odour  (see  footnote,  p.  159),  or  by  applying  the  reaction  with 
ammonium  sulphide  as  described  for  cyanogen — 

(NHJ2S2  +  HCy  =  (NH4)HS  +  NH4CyS 

Single  Cyanides. — These  are  binary  compounds,  analogous  to 
the  chlorides.  Of  these  the  cyanides  of  the  alkalies,  alkaline  earths, 
and  mercuric  cyanide  are  soluble  in  water  (barium  cyanide  with 
difficulty)  ;  all  others  are  insoluble.  The  aqueous  solutions  of 
these  cyanides  are  unstable,  undergoing  gradual  decomposition  ; 
chiefly  into  a  formate  and  free  ammonia — 

KCN  +  2H2O  =  K(CHO2)  +  NH3 

The  cyanides  of  the  alkali  metals  are  unchanged  when  strongly 
heated  alone,  while  those  of  the  heavy  metals  are  mostly  converted 
into  cyanogen  *  and  the  metal,  or  in  some  cases  into  nitrogen, 
carbon,  and  metal,  or  a  metallic  carbide.  Under  the  influence  of 
heat,  the  cyanides  of  the  alkalies  readily  take  up  oxygen  or  sulphur, 
from  compounds  capable  of  yielding  these  elements,  and  pass  into 
cyanate  or  thiocyanate.  Hence  potassium  cyanide  is  constantly 
employed  in  analysis  as  a  reducing  agent  in  blowpipe  reactions. 

The  alkali  cyanides,  and  some  of  the  insoluble  single  cyanides 
(e.g.  ZnCy2,  PbCy2)  are  readily  decomposed  by  dilute  mineral 
acids  ;  others  are  only  decomposed  with  difficulty  ;  in  either  case, 
with  liberation  of  hydrocyanic  acid.  They  are  all  decomposed 
by  strong  sulphuric  acid  (some  requiring  the  aid  of  heat),  with 
formation  of  metallic  sulphates  (see  p.  164).  Mercuric  cyanide  is 
decomposed  by  sulphuretted  hydrogen.  Cyanides  in  solution  may 
be  detected  (i)  by  the  precipitation  of  insoluble  single  cyanides  by 
double  decomposition  (not  very  characteristic)  ;  (2)  by  the  formation 
of  alkali  thiocyanate,  and  subsequently  obtaining  the  red  coloration 
with  ferric  chloride  ;  (3)  by  the  formation  of  Prussian  blue. 

(i)  Silver  nitrate,  added  to  a  solution  of  a  soluble  cyanide 
(other  than  mercuric  cyanide),  gives  a  white  precipitate  of  silver 
cyanide,  AgCy  ;  soluble  in  ammonia,  insoluble  in  nitric  acid,  and 
readily  soluble  in  potassium  cyanide,  giving  a  double  cyanide, 
KCy,AgCy. 

Silver  cyanide  closely  resembles  silver  chloride  ;  it  is,  however, 
readily  distinguished  by  the  fact  that,  when  boiled  with  hydrochloric 
acid,  it  is  decomposed,  with  evolution  of  hydrocyanic  acid ;  and 
also  that,  when  heated  alone,  it  gives  off  cyanogen  (detected  as 
above),  and  leaves  a  black  residue  of  silver  and  paracyanogen. 

*  And  usually  some  paracyanogen,  a  non-volatile  black  compound,  which 
is  a  polymeride  of  cyanogen  expressed  by  the  formula  (CN)jr. 


Hydrocyanic  Acid  and  Cyanides.  161 

(2)  Formation  of  Alkali  Thiocyanate. — Yellow  ammonium 
sulphide  is  added  to  the  solution  of  the  cyanide,  and  the  mixture 
gently  boiled  for  a  moment  or  two.    In  the  case  of  an  alkali  cyanide, 
the  following  change  takes  place  : — 

(NH4)2S2  +  KCy  =  KCyS  +  (NH4)2S 

With  mercuric  cyanide,  the  ammonium  sulphide  precipitates  black 
mercuric  sulphide  ;  thus — 

HgCy2  +  2(NH4)2S2  =  HgS  +  (NH4)2S  +  2NH4CyS 

A  solution  of  zinc  sulphate  is  then  added  (without  filtering,  in 
the  case  of  mercuric  cyanide),  until  the  excess  of  ammonium 
sulphide  is  removed  by  precipitation,  as  white  zinc  sulphide.  The 
mixture  is  then  filtered,  and  to  the  colourless  solution  a  drop  of 
ferric  chloride  is  added,  which  at  once  gives  the  deep  red  coloration 
due  to  ferric  thiocyanate.  (This  test  is  very  delicate.) 

(3)  Formation  of  Prussian  Blue.— To  the  solution  con- 
taining the  soluble  cyanide  (or  hydrocyanic  acid)  a  small  quantity  of 
sodium  or  potassium  hydroxide  is  added,  after  which  three  or  four 
drops  both  of  a  ferrous  and  ferric  salt  (ferrous  sulphate  and  ferric 
chloride).     The  visible  result  is  the  production  of  a  dirty  brown 
precipitate  of  the  mixed  ferrous  and  ferric  hydroxides.   The  invisible 
result  is  the  formation  of  ferrocyanide  in  the  solution.     The  ferrous 
sulphate,  interacting  with  the  cyanide  present,  forms  (i)  ferrous 
cyanide,  FeCy2,  which  in  its  turn  dissolves  in  the  excess  of  the 
cyanide,  forming  the  ferrocyanide  ;  thus,  with  potassium  cyanide — 

(1)  FeSO4  4-  2KCy  =  K2SO4  +  FeCy2 

(2)  FeCy2  +  4KCy  =  K4(FeCy6) 

On  the  addition  of  hydrochloric  acid  to  the  mixture,  the  pre- 
cipitated iron  hydroxides  are  dissolved,  and  the  ferric  chloride  thus 
formed  interacts  with  the  potassium  ferrocyanide  present,  giving 
rise  to  the  precipitation  of  ferric  ferrocyanide,  or,  in  cases  when 
the  amount  of  cyanide  is  very  small,  to  a  blue  or  green  coloration. 

Double  Cyanides. — Most  of  the  single  cyanides  are  soluble 
in  alkali  cyanides,  giving  rise  to  double  cyanides.  Some  of  these 
have  an  alkaline  reaction,  are  easily  decomposed  by  dilute  mineral 
acids  with  evolution  of  hydrocyanic  acid,  and  (like  the  single 
cyanides)  are  poisonous.  Others,  on  the  contrary,  have  a  neutral 
reaction,  do  not  give  off  hydrocyanic  acid  when  treated  with  dilute 
acids,  and  are  not  poisonous.  The  double  cyanides  are  therefore 
divided  into  two  classes,  namely  (i)  those  which  are  easily  decom- 
posed by  acids,  and  (2)  those  which  are  decomposed  with  difficulty. 

M 


1 62  Qualitative  Analysis. 

All  double  cyanides,  like  all  single  cyanides,  are  decomposed  by 
strong  sulphuric  acid,  and  also  by  fusion  with  potassium  nitrate 
(or  a  mixture  of  ammonium  nitrate  and  sulphate}. 

(i)  Easily  decomposed  Double  Cyanides.— From  an 
analytical  point  of  view,  these  may  all  be  regarded  as  mixtures 
of  the  single  cyanide  with  the  solvent  cyanide  (usually  potassium 
cyanide).  Thus,  the  double  cyanides  obtained  by  dissolving  the 
cyanides  of  zinc,  silver,  nickel,  in  potassium  cyanide,  are  expressed 
by  the  formulae  2KCy,ZnCy2,  KCy,AgCy,  2KCy,NiCy2.  Their 
solutions  contain  the  metals  as  the  positive  ions,  along  with  the 
anion  CN. 

Their  chemical  behaviour  is  practically  the  same  as  would  be 
shown  by  mixtures  of  potassium  cyanide  and  the  several  simple 
metallic  cyanides.  Thus,  when  such  double  cyanides  are  heated, 
the  alkali  cyanide  (which  is  stable  at  high  temperature)  remains, 
and  the  other  portion  of  the  molecule  behaves  exactly  as  the  single 
cyanide ;  eg.  the  silver  cyanide  evolves  cyanogen,  leaving  a  black 
residue  of  silver  and  paracyanogen. 

When  treated  with  acids,  the  same  thing  is  observed  ;  the 
alkali  cyanide  is  decomposed  with  evolution  of  hydrocyanic  acid 
(just  as  it  would  if  alone),  while  the  other  cyanide  in  the  molecule 
behaves  exactly  as  though  /'/  were  alone.  If  it  is  not  one  which  is 
decomposed  by  the  acid,  it  remains  ;  reprecipitated  of  course,  since 
the  solvent  cyanide  has  been  decomposed.  Thus,  the  double 
potassium  nickel  cyanide,  when  acted  upon  by  acids,  evolves  hydro- 
cyanic acid,  and  the  single  nickel  cyanide  is  reprecipitated — 

2KCy,NiCy2  +  2HC1  =  2KC1  +  2HCy  +  NiCy2 

In  the  case  of  the  double  potassium  zinc  cyanide,  not  only  is 
the  potassium  cyanide  decomposed,  but  also  the  zinc  cyanide  ; 
thus — 

2KCy,ZnCy2  +  4HC1  =  2KC1  +  ZnCl2  +  4HCy 

Some  of  these  double  cyanides  are  decomposed  by  sulphuretted 
hydrogen,  with  precipitation  of  the  sulphide  of  the  heavy  metal, 
e.g.  2KCy,HgCy2,  2KCy,CdCy2,  and  2KCy,ZnCy2.  In  other  cases, 
however,  sulphuretted  hydrogen  fails  to  precipitate  the  metals  as 
sulphides,  eg.  2KCy,CuCy2  or  6KCy,Cu2Cy2,  and  2KCy,NiCy2  (see 
also  pp.  60  and  83). 

All  the  double  cyanides  of  this  class  are  decomposed  by  being 
boiled  with  mercuric  oxide,  with  the  formation  of  mercuric  cyanide 
and  precipitation  of  the  oxide  or  hydroxide  of  the  heavy  metal — 


Ferrocyanides.  \  63 

The  detection  of  the  cyanogen  in  these  double  compounds  may 
be  accomplished  by  the  same  reactions  as  those  already  described 
for  single  cyanides.  To  recognise  the  heavy  metal  (as  in  the  case  of 
single  cyanides),  the  compound  must  be  completely  decomposed  by 
one  of  the  various  methods  given,  either  by  boiling  with  dilute  hydro- 
chloric acid,  or  by  strong  sulphuric  acid,  or  by  precipitation  with  mer- 
curic oxide,  or,  where  available,  by  means  of  sulphuretted  hydrogen. 

(2)  Difficultly  decomposed  Double  Cyanides.— These 
compounds  are  not  decomposed  with  liberation  of  hydrocyanic 
acid  by  dilute  mineral  acids.  They  may  be  regarded  as  containing 
a  complex  metallo-cyanogen  radical,  represented  by  the  general 
formula  RCy6 ;  and  in  the  case  of  those  which  are  of  most  im- 
portance in  analysis,  R  stands  for  Fe  or  Co,  as  in  the  following  : — 

Potassium  ferrocyanide,  K4(FeCy6) 

Copper  ferrocyanide,  Cu2(FeCy6) 

Potassium  ferrous  ferrocyanide,  K2Fe(FeCyG) 

Ferric  ferrocyanide  {Prussian  blue\  Fe4(FeCy6)3 

Potassium  ferricyanide,  Ks(FeCy6) 

Ferrous  ferricyanide  (Tumbvlfs  blue),  Fe3(FeCy6)2 

Potassium  cobalticyanide,  K3(CoCy6) 

The  reactions  which  potassium  ferro  and  ferri  cyanides  give 
with  ferric  and  ferrous  salts  have  already  been  studied  in  connection 
with  the  special  tests  for  iron.  The  two  salts  of  iron  are  therefore 
used  for  the  detection  of  soluble  ferro  and  ferri  cyanides. 

(a)  Ferrocyanides.— The  alkali  ferrocyanides  alone  are  easily 
soluble  in  water.  By  double  decomposition,  therefore,  with  potas- 
sium ferrocyanide  and  a  metallic  salt,  the  insoluble  ferrocyanides 
are  precipitated  ;  eg.-— 

K4(FeCy6)  +  2CuCl2  =  4KC1  +  Cu2(FeCy6) 
When   treated  with    cold   dilute   acids,  the  ferrocyanides  are 
decomposed,  not  with  evolution  of  hydrocyanic  acid,  but  with  the 
formation  of  hydroferrocyanic  acid,  the  complex  cyanogen  group 
remaining  intact— 

K4(FeCy6)  +  4HC1  =  4KC1  +  H4(FeCy6) 

When  boiled  with  dilute  sulphuric  acid,  the  alkali  ferro  (and 
"ferri")  cyanides  are  partially  decomposed,  with  elimination  of 
a  portion  of  the  cyanogen  as  hydrocyanic  acid.  (This  is  the  usual 
method  for  preparing  aqueous  hydrocyanic  acid)— 

2K4(FeCyc;  +  3H2SO4  -  K2SO4  +  K2Fe(FeCyc)  +  6HCy 


164  Qualitative  Analysis. 

When  heated  with  strong  sulphuric  acid  (in  common  with  all 
cyanides)  the  cyanogen  radical  is  completely  decomposed,  the 
nitrogen  it  contains  being  converted  into  ammonia— 

K4FeCy6+6H2SO4+6H2O  =  2K2SO4  +  FeSO4  +  3(NH4)2SO4  +  6CO 
Insoluble  ferrocyanides    are  decomposed   by  treatment  with 
caustic  alkalies  and  alkali  carbonates ;  thus,  with  Prussian  blue 
(ferric  ferrocyanide) — 

Fe4(FeCy6)3  +  i2NaHO  =  2Fe2(HO)6  +  3Na4(FeCy6) 

By  removing  the  precipitated  metallic  hydroxide  (in  this  case 
ferric  hydroxide)  by  filtration,  the  presence  of  the  ferrocyanide  in 
the  solution  can  be  detected  by  the  addition  of  ferric  chloride. 

Under  the  influence  of  oxidising  agents  (e.g.  chlorine  or  bromine 
water)  potassium  ferrocyanide  is  converted  into  ferricyanide — 

K4(FeCy6)  +  Cl  -  KC1  +  K3(FeCyG)  * 

(b}  Perricyanides.— The  alkali  salts  only  are  easily  soluble 
in  water.  Acids  and  alkalies  act  upon  the  ferricyanides  as  upon 
ferrocyanides ;  eg.  the  insoluble  compounds,  when  boiled  with 
alkaline  hydroxides,  yield  the  soluble  alkali  ferricyanide,  pre- 
cipitating the  metal  as  hydroxide  ;  thus,  with  ferrous  ferricyanide 
(Turnbull's  blue)— 

Fe3(FeCy6)2  +  6NaHO  -  2Na3(FeCy6)  +  3Fe(HO)2 

Soluble  ferricyanides  are  recognised  by  the  formation  of  a  blue 
precipitate  with  ferrous  sulphate  (see  Iron).  Insoluble  ferricyanides 
(in  common  with  all  other  double  cyanides)  are  examined  for  the 
metals  they  contain  by  first  decomposing  the  compound  ;  either  by 
boiling  with  alkali  hydroxide  as  shown  above,  or  by  boiling  with 
strong  sulphuric  acid  ;  or,  lastly,  by  fusion  with  a  mixture  of 
ammonium  nitrate  and  sulphate,  whereby  the  metals  are  converted 
into  sulphates. 

Cyanic  Acid  and  Cyauates. 

Cyanic  Acid,  HCyO. — This  compound  cannot  exist  in  con- 
tact with  water,  being  broken  up  into  carbon  dioxide  and  ammonia  ; 

*  Although  for  many  reasons  it  is  convenient  to  regard  these  compounds 
as  containing  the  complex  radical  (FeCy6),  this  view  of  their  constitution 
somewhat  obscures  some  of  their  reactions.  The  oxidation  of  potassium 
ferrocyanide  to  ferricyanide  is  a  case  in  point.  If  these  two  compounds  be 
represented  as  double  compounds  of  "  ferrous  "  and  "  ferric  "  cyanides  respec- 
tively, then  the  oxidising  action  of  chlorine  in  converting  the  one  into  the  other 
is  at  once  apparent — 

4KCy,Fe"Cy2  +  Cl  =  KC1  +  3KCy,Fe'"Cy3 


Thiocyanic  Acid  and  Thiocyanates.  165 

therefore,  when  cyanates  are  acted  upon  by  acids,  while  the  vapour 
of  cyanic  acid  escapes,  and  is  recognised  by  its  powerful  and 
pungent  or  acrid  smell,  the  greater  part  of  it  is  decomposed  by  the 
water  present ;  thus — 

2K(CN)0  +  2H2S04  +  2H20  =  K2S04  +  (NH4)2  SO4  +  2CO2 

Cyanates  of  the  alkalies  and  alkaline  earths  are  soluble  in  water. 
Insoluble  cyanates  are  therefore  precipitated  when  salts  of  most 
of  the  heavy  metals  are  added  to  a  solution  of  potassium  cyanate. 

Ammonium  cyanate,  NH4CNO,is  isomeric  with  urea,(NH2)2CO, 
and  when  an  aqueous  solution  of  the  former  is  heated,  it  passes 
into  the  latter,  by  simple  rearrangement  of  the  atoms  already  in 
the  molecule  without  any  addition  or  subtraction. 

Cyanates  are  at  once  distinguished  from  cyanides,  by  the  fact 
that  they  do  not  evolve  hydrocyanic  acid  on  treatment  with  dilute 
acids.  They  are  identified  by  the  evolution  of  carbon  dioxide, 
with  simultaneous  formation  of  ammonium  salt. 

Commercial  potassium  cyanide  usually  contains  potassium 
cyanate,  which  is  detected  by  this  reaction. 

Thiocyanic  Acid  and  Thiocyanates. 

Thiocyanates  are  the  sulphur  analogues  of  the  cyanates. 
The  acid  itself,  HCyS  (like  cyanic  acid),'  is  unstable  in  aqueous 
solution.  The  chief  insoluble  thiocyanates  are  those  of  silver,  lead, 
mercury,  and  copper.  On  the  addition  of  solutions  of  these  metals 
to  a  solution  of  potassium  thiocyanate,  the  various  compounds  are 
precipitated.  Silver  thiocyanate,  AgCyS,  is  obtained  as  a  white 
precipitate  resembling  the  chloride  in  appearance.  It  dissolves  in 
ammonia,  but  less  easily  than  either  the  cyanide  or  chloride,  and  is 
insoluble  in  dilute  nitric  acid. 

With  ferric  salts,  potassium  or  ammonium  thiocyanates  give  a 
red  colour,  which  is  not  destroyed  by  hydrochloric  acid.  The 
formation  of  this  colour  is  employed  both  as  a  test  for  "  ferric  "  iron 
or  for  thiocyanates.  The  addition  of  mercuric  chloride  destroys 
the  colour. 

When  heated,  the  thiocyanates  are  decomposed.  The  potassium 
salt  may  be  heated  to  fusion  in  the  absence  of  air,  but  in  contact 
with  air  it  is  converted  into  sulphate  and  cyanate,  with  evolution  of 
sulphur  dioxide.  The  mercurous  and  mercuric  salts,  when  heated 
in  contact  with  the  air,  take  fire  and  burn,  at  the  same  time 
swelling  up  in  a  most  remarkable  manner,*  leaving  a  straw-coloured 
voluminous  ash,  which  is  many  times  as  bulky  as  the  original 
compound.  At  the  same  time  sulphur  dioxide,  mercury  vapour, 
nitrogen,  and  cyanogen  are  evolved. 

*  The  thiocyanates  of  mercury  are  the  materials  of  which  the  so-called 
"  Pharaoh's  serpents  "  are  made. 


1 66  Qualitative  Analysis. 

Silicon. 

The  properties  of  the  element  silicon,  in  the  uncombined  state, 
are  not  made  use  of  in  analysis,  as  in  no  analytical  reactions  is  this 
element  liberated. 

Silicic  Acid  and  Silicates. 

Silicic  acid,  H2SiO3,  is  obtained,  when  soluble  silicates  are 
decomposed  by  acids,  as  a  white  gelatinous  substance,  slightly 
soluble  in  water,  and  still  a  little  more  soluble  in  acids — 

Na2SiO3  +  2HC1  =  H2SiO3  +  2NaCl 

If  the  decomposition  of  the  soluble  silicate  be  made  to  take  place 
in  the  presence  of  an  excess  of  acid— that  is,  if  the  silicate  be  added 
to  the  acid  (instead  of  adding  acid  to  the  silicate) — the  silicic  acid 
which  is  produced  is  not  precipitated  at  once,  but  remains  in  solu- 
tion, and  may  be  separated  from  the  sodium  chloride  and  the  excess 
of  hydrochloric  acid  by  dialysis.  After  a  short  time,  however,  the 
silicic  acid  separates  out  as  a  transparent  gelatinous  mass.* 

Silicic  acid  is  also  precipitated  from  a  solution  of  an  alkali 
silicate  by  the  addition  of  ammonium  carbonate  (ammonia  does 
not  cause  any  precipitate).  The  action  may  be  regarded  as  taking 
place  in  two  stages,  ammonium  silicate  being  first  formed  by 
double  decomposition,  and  immediately  breaking  up  into  silicic 
acid  and  ammonia  ;  thus — 

Na2Si03  +  (NH4)2C03  =  Na2CO3  +  (NH4)2SiO3 
(NH4)2SiO3  =  2NH3  +  H2SiO3 

Silicic  acid  is  a  very  feeble  acid,  and  when  heated  to  130°  C. 
it  parts  with  a  molecule  of  water  and  is  converted  into  silicon 
dioxide  (silica],  SiO2,  a  compound  which  is  insoluble  in  water  and 
in  acids.t 

Silica,  SiO2,  occurs  in  a  more  or  less  pure  state  in  nature,  in  the 
form  of  quartz,  flint,  agate,  sand,  etc.  When  prepared  artificially 
by  heating  the  hydrated  compound,  it  is  a  white  amorphous  powder. 
It  is  extremely  stable,  and  capable  of  standing  a  very  high  tempera- 
ture. It  is  insoluble  in  water  and  all  acids  except  hydrofluoric 
acid.  It  dissolves  in  caustic  alkalies  ;  and,  when  in  the  amorphous 

*  The  acid  which  is  thus  obtained  in  solution  is  believed  to  be  the  less 
stable  tetrabasic  silicic  acid,  H4SiO4,  or  Si(HO)4,  or  SiOj.aHjO.  This  com- 
pound, by  giving  up  one  molecule  of  water,  passes  into  the  dibasic  acid,  H2SiO3, 
or  SiO2,H2O. 

f  Silicic  acid  is  sometimes  spoken  of  as  soluble  silica,  and  silicon  dioxide  as 
insoluble  silica. 


Silicic  Acid  and  Silicates.  167 

state,  in  boiling  carbonates  of  the  alkalies,  yielding  in  all  cases  the 
soluble  alkali  silicates. 

The  only  silicates  which  are  soluble  in  water  are  the  alkali 
silicates  (known  as  water  glass,  or  soluble  glass).  The  natural 
silicates  form  a  numerous  and  complex  class  of  minerals.  They 
may  be  regarded  as  compounds  of  metallic  oxides  with  silicon 
dioxide  ;  thus,  felspar,  Al2O3JK2O,6SiO2. 

The  silicates,  which  are  insoluble  in  water,  may  be  classified  for 
analytical  purposes  into  (i)  those  which  are  decomposed  by  acids 
(other  than  hydrofluoric  acid) ;  and  (2)  those  which  are  unattacked, 
which  comprises  by  far  the  larger  class. 

The  presence  of  silica  in  these  compounds  is  detected  by  means 
of  their  behaviour  when  heated  with  microcosmic  salt.  A  clear 
bead  of  this  salt  is  made  upon  a  loop  of  platinum  wire,  and  a  few 
small  particles  of  the  powdered  mineral  are  heated  in  it.  The 
metallic  oxides  are  dissolved  by  the  fused  microcosmic  salt,  but  not 
the  silicon  dioxide,  particles  of  which  (the  skeletonic  remains  of  the 
mineral)  remain  floating  about  in  the  molten  bead.* 

Silica  may  also  be  detected  by  the  formation  of  silicon  fluoride, 
when  a  silicate  is  acted  upon  by  hydrofluoric  acid.  The  test  is 
made  as  described  under  fluorine  (p.  134).!  In  this  case,  however, 
the  drop  of  water  which  is  suspended  in  the  gas  in  order  to  detect 
the  silicon  fluoride,  should  be  held  on  a  loop  of  platinum  wire,  as  the 
action  of  the  hydrofluoric  acid  upon  the  glass  rod  might  be  mis- 
taken for  a  deposition  of  silica.  When  applying  this  test  for  the 
detection  of  fluorine,  this  is  obviously  no  disadvantage. 

(i)  Silicates  which  are  decomposed  by  acid  may  have 
their  silica  removed  by  treatment  with  hydrochloric  acid,  just  as 
in  the  case  of  the  alkali  silicates  above  mentioned.  In  this  instance, 
however,  the  mineral,  in  as  finely  powdered  a  condition  as  possible, 
is  digested  with  strong  hydrochloric  acid  at  a  gentle  heat,  when  gela- 
tinous silicic  acid  separates,  and  the  powdered  mineral  gradually 
dissolves.  In  order  to  render  the  separation  of  the  silica  complete, 
the  mixture  is  next  evaporated  to  dryness  (preferably  on  a  steam- 
bath),  and  in  order  to  ensure  the  entire  conversion  of  the  silicic  acid 
into  silica,  it  is  afterwards  gently  heated  over  a  flame  for  a  short  time  : 
(excessive  heating  is  to  be  avoided,  as  certain  metallic  oxides  are 
thereby  rendered  very  difficult  of  subsequent  solution  in  acid).  The 
residue  is  then  treated  with  a  small  quantity  of  strong  hydrochloric 

*  This  preliminary  test,  however,  is  not  absolutely  reliable.  Some  in- 
soluble silicates  are  known  which  do  not  leave  this  residue,  and  there  are 
certain  natural  phosphates  which  are  not  dissolved  by  fused  microcosmic  salt, 
but  remain  floating  in  the  bead,  and  might  be  mistaken  for  silica. 

t  Except  that  a  platinum  crucible  must  replace  the  test-tube. 


1  68  Qualitative  Analysis. 

acid  ;  water  is  added,  and  the  solution  containing  the  metals  present 
as  chlorides  is  separated  by  nitration  from  the  insoluble  residue  of 
silica. 

(2)  Silicates  which  are  unattacked  by  acids  are  decom- 
posed by  fusion  with  alkali  carbonates.  The  finely  powdered 
silicate  is  mixed  with  several  times  its  weight  of  fusion  mixture,  and 
the  mixture  heated  in  a  platinum  crucible.  As  the  mass  melts, 
effervescence  takes  place,  due  to  the  escape  of  carbon  dioxide  ; 
thus,  using  a  general  formula  for  a  simple  silicate  *  — 


The  fusion  is  continued  until  effervescence  ceases,  the  temperature 
being  raised  towards  the  end  of  the  operation.  The  residue,  after 
cooling,  is  extracted  with  water,  which  dissolves  the  alkali  silicate 
(and  excess  of  carbonate),  leaving  the  metallic  oxide  (or  carbonate). 
Hydrochloric  acid  is  then  gradually  added,  which  causes  the  pre- 
cipitation of  silicic  acid,  and  at  the  same  time  dissolves  the  metallic 
oxides.  The  mixture  is  then  evaporated,  and  treated  in  the  manner 
described  above. 

The  acid  radicals  of  any  other  compounds  present  in  the  silicate 
(eg.  sulphates,  phosphates,  borates,  arsenates,  chlorides)  are  now 
present  as  their  sodium  or  potassium  salts,  and  may  be  detected 
by  their  several  special  tests  in  the  ordinary  way. 

This  method  of  fusion  with  alkali  carbonates  is  obviously  inad- 
missible in  the  case  of  natural  silicates  which  are  suspected  of 
containing  the  alkali  metals,  and  which  are  insoluble  in  hydrochloric 
acid.  In  this  case,  one  of  the  following  plans  may  be  employed:  — 

(a)  Heating  with  Barium  Oxide.—  A  small  quantity  of  the 
finely  powdered  mineral  is  intimately  mixed  with  three  or  four 
times  its  weight  of  barium  oxide,  and  strongly  heated  in  a  platinum 
crucible  by  the  blowpipe,  the  heating  being  continued  for  about 
twenty  minutes.  The  mass  is  then  dissolved  in  dilute  hydrochloric 
acid.  The  solution  so  obtained  is  made  alkaline  with  ammonia  (a 
precipitate  is  produced,  which  need  not  be  removed),  and  ammo- 
nium carbonate  is  added.  The  mixture  is  then  filtered,  and  the 
filtrate  evaporated  to  dryness  and  heated,  and  examined  for  the 
alkalies  in  the  usual  way  (p.  25). 

(/;)  Heating  with  Ammonium  Chloride  and  Calcium 
Carbonate.  —  A  small  quantity  of  the  powdered  silicate  is  mixed 
with  about  an  equal  weight  of  ammonium  chloride  and  about  eight 

*  The  possibility  of  this  reaction  depends  upon  the  fact  that,  although 
silicic  acid  is  an  unstable  acid,  silicon  dioxide  and  the  alkaline  silicates  are 
extremely  stable  even  at  high  temperatures. 


Boric  Acid  and  B orates.  169 

times  its  weight  of  calcium  carbonate  (precipitated),  and  the  mix- 
ture gently  heated  in  a  platinum  crucible  until  no  more  fumes  of 
ammonium  chloride  are  given  off.  It  is  then  strongly  heated  with 
the  blowpipe  for  about  fifteen  minutes,  after  which  the  mass  is 
treated  with  water.  The  alkalies,  in  the  form  of  chlorides,  together 
with  a  small  quantity  of  calcium  chloride,  pass  into  solution,  and 
are  separated  by  filtration.  The  calcium  is  removed  by  pre- 
cipitation with  ammonium  carbonate,  and  the  filtrate  evaporated 
and  examined  for  alkalies. 

(c)  Decomposing  the  Silicate  with  Hydrofluoric  Acid. — 
This  may  be  accomplished  by  treating  the  powdered  mineral  with 
aqueous  hydrofluoric  acid  in  a  platinum  crucible,  gently  evaporating 
the  liquid  (in  a  draught  cupboard)  to  dryness,  adding  fresh  acid 
and  evaporating  again,  continuing  the  operation  until  the  residue  is 
entirely  soluble  in  hydrochloric  acid.  Or  the  mineral  may  be 
moistened  with  a  little  strong  ammonia  in  a  platinum  crucible,  and 
the  whole  exposed  to  the  action  of  gaseous  hydrofluoric  acid 
(generated  from  fluorspar  and  sulphuric  acid)  in  a  leaden  pot,  which 
can  be  covered  with  a  leaden  lid.  The  action  is  allowed  to  proceed 
for  at  least  twenty-four  hours.  The  crucible  is  then  removed,  and 
the  ammonium  fluoride  expelled  by  gently  heating  in  a  draught 
cupboard. 

Boron. 

The  properties  of  the  free  element  are  not  employed  in  analysis. 

Boric  Acid  and  Borates. 

Boric  acid,  H3BO3,  is  a  white  crystalline  solid,  sparingly 
soluble  in  cold,  but  more  readily  soluble  in  hot,  water.  It  is 
deposited  from  its  solutions  in  the  form  of  pearly  white  scales. 

It  is  also  soluble  in  alcohol,  and  when  either  the  aqueous  or 
alcoholic  solution  is  boiled,  the  acid  vaporises  along  with  the  solvent. 

When  heated,  boric  acid  (ortho),  H3BO3,  loses  water,  passing  first 
into  metaboric  acid,  H2B2O4,  and  finally  into  pyroboric  acid,  H2B4O7. 

The  borates  of  the  alkalies  are  readily  soluble  in  water ;  most 
other  borates  are  insoluble.  A  few  (e.g.  magnesium  borate)  are 
difficultly  soluble.  The  most  familiar  salts  are  those  of  pyroboric 
acid,  e.g.  ordinary  borax,  Na2B4O7. 

The  insoluble  borates  obtained  by  precipitation  are  not  charac- 
teristic, and  are  not  used  in  analysis  for  the  detection  of  borates. 

All  borates  are  decomposed  by  mineral  acids,  with  liberation  of 
orthoboric  acid  ;  thus — 

Na2B4O7  +  2HC1  4-  5H,O  =  2NaCl  +  4H3BO3 


170  Qualitative  Analysis. 

Reaction  with  Turmeric.— Boric  acid  produces  upon  tur- 
meric paper  a  characteristic  red-brown  stain.  This  coloration  is 
distinguished  from  that  produced  by  alkalies  (which  it  closely  re- 
sembles in  appearance)  by  the  fact  that  when  touched  with  an  alkali 
the  brown  colour  is  changed  to  a  greenish-black,  but  is  restored  to 
its  original  tint  by  dilute  acids  (HC1  or  H2SO4).  The  borate  is 
moistened  with  hydrochloric  or  sulphuric  acid  in  order  to  liberate 
the  boric  acid,  and  a  drop  or  two  of  the  liquid  is  poured  upon  the 
turmeric  paper. 

Flame  Reactions.  —  Volatile  boron  compounds  impart  a 
characteristic  green  colour  to  a  non-luminous  flame  ;  thus,  when 
an  alcoholic  solution  of  boric  acid  is  boiled,  and  the  alcohol  vapour 
inflamed,  the  green  colour  due  to  the  volatilised  boric  acid  is 
apparent.  The  test  is  made  in  the  following  manner  :— 

The  borate  (borax)  is  moistened  with  a  little  strong  sulphuric 
acid  in  a  test-tube  or  small  flask,  and  alcohol  is  added.  The  test- 
tube  is  closed  with  a  cork  carrying  a  short  straight  glass  tube. 
The  contents  of  the  tube  are  then  heated,  and  as  the  alcohol  boils 
off  it  is  inflamed  at  the  exit  tube,  when  the  green  colour  of  the 
flame  is  observed. 

Since  ethyl  chloride  (which  is  liable  to  be  formed  if  a  metallic 
chloride  is  treated  in  the  same  manner)  gives  a  flame  which  also 
has  a  green  edge  or  fringe  (albeit,  if  this  compound  were  formed, 
the  alcohol  flame  would  at  once  be  rendered  luminous,  like  an 
ordinary  gas  flame),  the  following  test  may  also  be  applied  as  a 
confirmation  : — 

A  small  quantity  of  the  powdered  borate  (or  boric  acid)  is 
mixed  with  about  its  own  weight  of  powdered  calcium  fluoride 
and  three  or  four  times  its  weight  of  hydrogen  potassium  sulphate. 
This  mixture  is  moistened  with  the  least  drop  of  water,  and  a  little 
of  the  paste  is  introduced  into  a  Bunsen  flame  upon  a  loop  of 
platinum  wire.  By  the  action  of  the  acid  sulphate  upon  the 
fluoride,  hydrofluoric  acid  is  formed  ;  and  this  in  the  presence 
of  the  borate  gives  boron  fluoride,  BF3,  which  causes  a  green 
coloration  of  the  flame  : — 

(1)  CaF2  +  2HKSO4  =  CaSO4  +  K2SO4  +  2HF 

(2)  B203  +  6HF  =  3H20  +  2BF3 

In  many  cases  it  is  sufficient  to  mix  the  borate  with  the  fluoride, 
and  moisten  the  mixture  with  a  drop  of  strong  sulphuric  acid  and 
bring  this  into  the  flame  in  the  same  way.  The  presence  of  copper 
salts  masks  the  reaction. 


CHAPTER  XV. 
SYSTEMATIC  DETECTION  OF  THE  ACIDS. 

THE  acids  are  classified  into  three  main  groups  based  upon  the 
solubility  of  their  barium  and  silver  salts. 

Group  I. — Acids  whose  barium  salts  are  precipitated  from 
neutral  solutions  by  barium  chloride. 

(a)  Whose  barium  salts  are  insoluble  in  dilute  hydrochloric  acid- 
Sulphuric  acid,  H2SO4  Hydrofluosilic  acid,  H2SiF6* 

(£)  Whose  barium  salts  are  soluble  in  hydrochloric  acid — 

Carbonic  acid,          H2CO3  Boric  acid,              H3BO3 

Sulphurous,,             H2SO3  lodic     „                  HIO3 

Thiosulphuric  acid,  H2S2O3  Hydrofluoric  acid,  HF 

Silicic                 „      H2Si03  Oxalic              „      H2(C2O4)t 

Chromic             „      H2CrO4  Tartaric           „       H2(C4H4O6)t 

Phosphoric        „      H3PO4  Citric                „      H3(C6H5O7)t 

Group  II. — Acids  whose  silver  salts  are  precipitated  by  silver 
nitrate  from  solutions  acidified  with  nitric  acid — 

Hydrochloric  acid,  HCl  Sulphuretted  hydrogen,  H2S 

Hydrobromic    „       HBr  Thiocyanic  acid,  HCyS 

Hydriodic         „       HI  Ferrocyanic  „  H4FeCy6 

Hydrocyanic    „       HCy  Ferricyanic   „  H3FeCy6 

Group  III. — Acids  whose  barium  salts  are  soluble  in  water, 
and  whose  silver  salts  are  not  precipitated  in  a  nitric  acid  solution, 
namely — 

Nitric  acid,    HNO3 

Nitrous  „       HNO2  Cyanic  acid,  HCyO 

*  Soluble  in  strong  hydrochloric  acid. 

f  In  the  case  of  these  three  acids,  the  neutral  solution  must  not  contain 
appreciable  quantities  of  ammoniacal  salts,  or  soluble  double  compounds  are 
formed  on  addition  of  barium  chloride,  which  may  entirely  prevent  the  forma- 
tion of  a  precipitate. 


172  Qualitative  Analysis. 

Chloric  acid,       HCIO3  Formic  acid,  H(CHO2) 

Perchloric  acid,  HC1O4  Acetic     „      H(C2H3O2) 

As  already  mentioned  (p.  120),  these  general  reagents,  barium 
chloride  and  silver  nitrate,  are  not  employed  to  effect  the  separa- 
tion of  groups  of  acids,  but  rather  as  indicators  of  the  presence  or 
absence  of  entire  groups. 

In  most  cases,  each  acid  is  individually  tested  for  in  separate 
portions  of  specially  prepared  solutions  (see  below). 

The  examination  for  acids  should  be  made  after  the  metals 
have  been  detected,  for  two  reasons  :  firstly,  because  during  the 
course  of  the  systematic  examination  for  the  metals,  the  presence 
or  absence  of  quite  a  number  of  acids  will  be  incidentally  ascertained ; 
and  secondly,  because  a  knowledge  of  what  metals  are  present 
will  be  a  guide  to  the  student  in  deciding  what  acids  may  and  what 
cannot  possibly  be  present.  For  example,  if  the  material  under 
examination  is  a  solution  having  an  acid  reaction,  and  it  is  found 
to  contain  silver  as  one  of  the  metals,  it  will  obviously  be  useless 
to  expect  to  find  any  of  the  acids  of  Group  II.  Or  if  the  substance 
is  a  solid  which  dissolves  easily  in  water,  and  calcium  is  found,  it 
is  equally  unnecessary  to  apply  tests  for  sulphuric,  phosphoric, 
carbonic,  silicic,  hydrofluoric  or  other  acids  whose  calcium  salts 
are  not  easily  soluble  in  water.  Or  if  a  solid,  soluble  in  water, 
were  found  to  contain  the  metal  lead,  the  acids  most  probably 
present  would  be  nitric  or  acetic,  as  the  lead  salts  of  these  two 
acids  are  the  only  common  soluble  lead  compounds.  A  student 
who,  in  spite  of  this,  goes  blindly  through  a  series  of  tests  for  acids 
which  cannot  be  present,  shows  at  once  that  he  has  no  intelligent 
appreciation  of  the  work  he  is  doing. 

It  will  be  evident  that,  in  order  to  make  the  fullest  use  of  the 
information  he  has  gained  concerning  the  metals  that  are  present, 
the  student  must  have  a  knowledge  of  the  solubilities  of  a  large 
number  of  compounds.  The  solubility  of  the  common  compounds 
of  the  metals  and  acids  has  been  mentioned  under  the  separate 
elements  treated  in  the  foregoing  chapters  ;  but  in  order  to  make 
the  information  more  accessible  for  reference,  a  condensed  epitome 
of  solubilities  of  the  common  compounds  of  the  various  metals  and 
acids  which  have  been  treated,  is  given  at  the  end  of  this  chapter 

(P-  177). 

As  stated  above,  several  acids  will  be  detected  during  the 
examination  for  metals  ;  thus — on  acidifying  with  hydrochloric 
acidtt\&  gently  warming  (for  separation  of  Group  I.),  the  presence 
of  carbonates,  sulphites,  sulphides,  cyanides,  or  nitrites  will  be 


Systematic  Detection  of  the  Acids.  173 

indicated  by  the  evolution  of  carbon  dioxide,  sulphur  dioxide, 
sulphuretted  hydrogen,  hydrocyanic  acid,  or  nitric  oxide  (appearing 
as  brown  fumes  of  the  peroxide)  respectively.* 

The  evolution  of  sulphur  dioxide  with  simultaneous  deposition 
of  sulphur  indicates  thiosulphates.  The  escape  of  chlorine  may 
denote  hypochlorites,  or  may  arise  from  the  action  of  the  hydro- 
chloric acid  (especially  if  not  very  dilute)  upon  chlorates,  iodates, 
chromates,  nitrates,  or  (in  the  case  of  solid  substances)  of  peroxides. 

When  the  substance  under  examination  is  a  solid,  these  indi- 
cations obtained  by  the  action  of  hydrochloric  acid  should  be 
followed  up  by  a  second  general  test,  namely — 

Treatment  with  Strong  Sulphuric  Acid. — On  gently  warm- 
ing a  little  of  the  solid  with  a  small  quantity  of  concentrated 
sulphuric  acid  in  a  test-tube,  a  number  of  important  indications 
may  be  obtained  ;  thus,  Cyanides,  chlorides,  fluorides,  and  nitrates 
evolve  their  respective  acids,  in  the  latter  case  accompanied  by 
brown  fumes.  Iodides  give  violet  vapours ;  bromides  give  the 
dark-brown  vapour  of  bromine,  condensing  to  dark  drops  on  the 
tube ;  chlorates  yield  the  deep  yellow  explosive  gas,  chlorine 
peroxide.  Formates  evolve  carbon  monoxide,  while  oxalates  give 
carbon  monoxide  and  dioxide  (the  residue  in  neither  case  is  black- 
ened). Tartrates  and  citrates  evolve  sulphur  dioxide  in  addition 
to  the  oxides  of  carbon,  at  the  same  time  leaving  a  black  residue. 
In  many  cases  these  indications  should  be  followed  up  at  once  by 
special  confirmatory  tests,  which  it  may  be  practicable  to  apply 
directly  to  the  substance  under  investigation;  e.g.  an  indication 
which  leads  to  the  suspicion  of  nitrates  being  present,  should  be 
confirmed  by  applying  the  test  with  copper,  or  the  ferrous  sulphate 
reaction.  If,  therefore,  this  test  with  sulphuric  acid  should  give 
a  purely  negative  result,  it  is  obvious  that  a  large  number  of  acids 
are  excluded.  The  chief  compounds  which  give  no  indications 
when  subjected  to  this  test  are  sulphates,  phosphates,  phosphites, 
arsenates,  silicates,  and  metallic  oxides  other  than  peroxides,  or 
such  as  behave  as  peroxides  ;  these  latter  evolve  oxygen. 

On  treatment  with  sulphuretted  hydrogen,  in  the  separation 
of  the  metals  of  Group  II.,  indications  will  have  been  obtained 
of  the  presence  of  chromates  and  iodates :  the  former  by  the 
change  of  colour  from  orange  to  green,  with  simultaneous  pre- 
cipitation of  sulphur  ;  the  latter  by  the  elimination  of  iodine,  which 

*  If  these  indications  of  the  decomposition  of  salts  of  these  acids  were  not 
carefully  noted  during  the  examination  for  metals,  the  reaction  should  be 
repeated  at  this  stage  of  the  analysis. 


174  Qualitative  Analysis. 

gives  a  dark  brown  colour  to  the  solution,  the  colour  gradually 
disappearing  as  the  excess  of  sulphuretted  hydrogen  converts  the 
iodine  into  hydriodic  acid. 

On  the  preparation  of  the  solution  for  separating  the  metals  of 
Group  III.,  the  presence  or  absence  of  phosphates  and  silicates 
will  have  been  ascertained,  and  also  of  such  organic  acids  (citric 
or  tartaric)  as  leave  a  charred  residue  when  heated. 

Besides  these,  indications  of  the  presence  of  several  acids  will 
have  been  obtained  during  the  performance  of  the  general  dry 
tests  for  the  metals  (p.  179). 

Preparation  of  the  Solution  for  the  Detection  of  Acids. 

Before  testing  for  the  acids,  it  is  advisable  (in  many  cases  it  is 
necessary]  that  they  should  be  present  in  the  solution  as  salts  of 
the  alkalies  (or  alkaline  earths)  ;  that  is  to  say,  the  metals  with 
which  the  various  acids  are  united  in  the  substance  under  analysis 
should  be  exchanged,  by  double  decomposition,  for  an  alkali  metal ; 
for  the  reason  that  the  presence  of  these  other  metals  would  in  many 
cases  mask  the  reactions  by  which  the  acids  are  to  be  detected.* 
To  accomplish  this,  the  solution  is  boiled,!  and  sodium  carbonate 
added  in  quantity  slightly  in  excess  of  that  required  to  effect 
complete  precipitation.  The  mixture  is  then  filtered,  and  the 
acids,  now  present  as  their  sodium  salts,  are  detected  in  the 
solution.  % 

*  Whether  or  not  this  would  be  the  case  will  usually  be  revealed  by  the 
result  of  the  examination  for  metals.  For  example,  suppose  the  metals  found 
in  a  solution  were  magnesium  and  copper,  then  the  only  acids  likely  to  be 
present  would  be  hydrochloric,  sulphuric,  and  nitric  (these  acids  forming  the 
common  soluble  salts  of  the  two  metals),  and  the  reactions  by  which  these 
acids  may  be  detected  can  be  made  equally  well  whatever  metal  they  may  be 
united  with.  Or  again,  suppose,  in  addition  to  these  two  metals,  silver  was 
found ;  this  would  reduce  the  acids  probably  present  to  nitric  and  sulphuric, 
and  the  test  for  the  latter  would  then  merely  require  the  substitution  of  barium 
nitrate  for  barium  chloride.  The  student  should  therefore  use  an  intelligent 
judgment  at  this  point,  and  not  proceed  by  a  blind  habit  to  prepare  the 
solution  by  the  removal  of  the  metals  with  sodium  carbonate,  whether  such  a 
step  be  necessary  or  not.  The  importance  of  carefully  and  systematically 
taking  notes  of  his  work  as  it  is  in  progress  cannot  be  too  strongly  urged. 
Such  notes  should  be  recorded  as  important  memoranda  for  his  own  guidance 
and  instruction,  and  not  as  a  mere  voucher  for  his  teacher  that  he  has  carried 
out  a  certain  prescribed  process,  or  made  a  certain  stereotyped  set  of  tests. 

t  In  cases  where  the  substance  under  analysis  is  insoluble,  the  product 
obtained  by  fusion  with  sodium  carbonate  is  extracted  with  water,  and  the 
solution  so  obtained  is  employed  for  the  detection  of  the  acids  (see  p.  184). 

J  This  method  for  separating  the  metals  from  their  acids  is  not,  however, 
of  universal  application.  Thus,  in  the  case  of  many  of  the  phosphates  held 
in  solution  by  acids,  the  action  of  the  sodium  carbonate  is  to  cause  the  pre- 
cipitation of  the  phosphates  themselves ;  in  such  instances,  therefore,  the  acid 
is  not  found  in  the  solution.  Since,  however,  phosphoric  acid  will  have  been 


Systematic  Detection  of  the  Acids.  175 

GENERAL  TESTS.— (i)  A  small  portion  of  this  solution  is  care- 
fully neutralised  by  first  adding  dilute  nitric  acid  drop  by  drop 
(the  mixture  being  heated  to  expel  carbon  dioxide),  until  the  liquid 
is  just  acid,  and  then  adding  a  drop  or  two  of  dilute  ammonia. 

This  neutral  solution  is  then  tested  by  the  addition  of  barium 
chloride.  If  no  precipitate  is  obtained,  the  acids  of  Group  I. 
(p.  171)  are  absent.  If  a  precipitate  is  formed  which  redissolves  on 
the  addition  of  hydrochloric  acid,  the  two  acids  sulphuric  and 
hydrofluosilicic  are  excluded. 

(2)  A  second  small  portion  of  the  solution  is  acidified  with 
nitric  acid,  and  tested  with  silver  nitrate.  A  negative  result  proves 
the  absence  of  the  acids  of  Group  II. 

If  these  general  tests  show  that  acids  of  both  Groups  I.  and  II. 
are  present,  special  tests  must  then  be  applied  in  separate  portions 
of  the  solution,  for  such  acids  as,  from  information  already  gained, 
are  considered  likely  to  be  in  the  substance  under  analysis,  and 
which  have  not  been  definitely  discovered  during  the  course  of  the 
examination.  The  tests  may  be  made  in  accordance  with  the 
following  outline  scheme,  in  which  most  of  those  acids  which  must 
certainly  have  been  detected  in  the  earlier  stages  are  not  again 
mentioned. 

A.  In  portions  of  the  solution  acidulated  with  hydrochloric  acid. 

Sulphuric  Acid. — Barium  chloride  precipitates  white  barium 
sulphate  :  not  dissolved  by  the  addition  of  strong  hydrochloric 
acid,  and  boiling  the  liquid. 

Hydrofluosilic  Acid. — Barium  chloride  gives  white  barium 
silicofluoride  :  soluble  in  strong  hydrochloric  acid. 

Ammonia  precipitates  gelatinous  silicic  acid. 

Silicic  Acid.— Ammonium  carbonate  (but  not  ammonia)  pre- 
cipitates silicic  acid. 

"  Perro  "  and  "  ferri  "-cyanic  acids  *  give  their  respective 
reactions  with  iron  salts. 

Thiocyanic  acid  *  gives  the  red  coloration  with  ferric  salts. 

discovered  during  the  examination  for  the  metals,  this  fact  is  of  little 
consequence. 

On  the  other  band,  sodium  carbonate  fails  to  separate  the  metals  from 
cyanides  soluble  in  water  (mercuric  cyanide)  or  In  potassium  cyanide  (i.e. 
soluble  double  cyanides,  such  as  those  of  silver,  copper,  zinc,  etc.),  and  also 
from  double  tartrates  and  citrates,  etc.  In  such  cases  as  these,  the  metals 
may  be  separated  by  means  of  sulphuretted  hydrogen  or  ammonium  sulphide 
(see  Reactions  for  cyanides,  p.  161). 

It  will  be  obvious  that  carbonates  must  be  detected  at  an  earlier  stage  in  the 
analysis. 

*  If  the  general  test,  No.  2  (see  above),  gave  a  negative  result,  these  acids 
will  have  been  proved  to  be  absent. 


176  Qualitative  Analysis. 

Arsenic  Acid. — If  arsenic  has  been  found  among  the  metals 
and  its  state  of  oxidation  (i.e.  whether  present  as  an  arsenite  or 
arsenate)  has  not  been  determined,  the  following  test  may  be 
applied  :  Ammonium  chloride,  ammonia,  and  magnesium  sulphate 
are  added — a  white  precipitate  may  be  due  to  either  a  phosphate 
or  arsenate.  If  phosphoric  acid  has  been  proved  to  be  absent  (by 
the  molybdate  reaction),  it  must  be  the  arsenate.  In  either  case 
it  should  be  filtered,  and  after  being  washed  free  from  ammonium 
chloride,  it  is  dissolved  in  a  little  nitric  acid,  and  silver  nitrate 
added  (or  a  drop  or  two  of  silver  nitrate  may  be  poured  upon  the 
washed  precipitate  in  the  funnel).  A  brown  precipitate  indicates 
an  arsenate. 

B.  In  portions  acidulated  with  nitric  acid. 
Hydrochloric,  Hydrocyanic,  and  Thiocyanic  Acids. 

Silver  nitrate  gives  a  white  precipitate. 

Hydrobromic  and  Hydriodic  Acids. — Silver  nitrate  gives 
a  yellowish  precipitate. 

(For  the  methods  of  discriminating  between  these,  see  pp.  127 
and  128.) 

C.  In  portions  acidulated  with  acetic  acid. 

Oxalic  Acid. — Calcium  sulphate  precipitates  calcium  oxalate  : 
readily  soluble  in  hydrochloric  acid. 

Hydrofluoric  Acid. — Calcium  sulphate  or  chloride  precipi- 
tates calcium  fluoride  :  only  slightly  soluble  in  hydrochloric  acid. 

(Chromic,  phosphoric,  and  arsenic  acids  may  also  be  looked 
for  in  portions  of  this  solution.  These  acids,  however,  will  have 
been  detected  at  an  earlier  stage  of  the  analysis.) 

D.  In  portions  rendered  neutral. 

Tartaric  and  Citric  Acids.— Calcium  chloride  gives  a  white 
precipitate  of  calcium  tartrate  in  the  cold,  and  of  calcium  citrate 
on  boiling  (see  Special  reactions,  p.  157). 

Formic  and  Acetic  Acids. — Ferric  chloride  produces  a  red- 
coloured  solution,  browner  in  colour  than  that  given  by  thio- 
cyanates.  The  colour  is  destroyed  by  hydrochloric  acid  (see 
Special  reactions). 


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CHAPTER  XVI. 

PRELIMINARY  TESTS   AND  OPERATIONS   IN   A   SYSTEMATIC 
ANALYSIS. 

A.  When  the  Substance  under  Examination  is  a  Liquid. 

— Before  proceeding  to  the  separation  of  the  groups,  the  liquid 
should  be  carefully  tested  with  litmus  paper,  in  order  to  ascertain 
whether  it  is  neutral,  acid,  or  alkaline. 

(1)  If  neutral^  the  liquid  might  be  simply  water.     To  ascertain 
whether  or  not  it  contains  anything  in  solution,  a  little  of  it  should 
be  carefully  evaporated  to  dryness.     This  should  be  done  upon  a 
watch-glass,  which  is  heated  upon  a  steam-bath.     (The  glass  may 
be  placed  over  the  mouth  of  a  small  beaker,  in  which  a  little  water 
is  being  gently  boiled.)    If  there  is  no  residue  left  in  the  watch-glass 
when  the  liquid  has  entirely  evaporated,  it  contained  no  salts  in 
solution,  and  was  therefore  simply  water. 

[Unless  evaporated  gently  as  described,  certain  compounds 
which  could  be  present  might  volatilise  entirely,  and  in  this  way  be 
altogether  overlooked  ;  e.g.  ammonium  nitrate  would  go  off  as 
steam  and  nitrous  oxide,  if  its  aqueous  solution  were  merely  boiled 
to  dryness.] 

(2)  If  acid,  the  solution  may  contain  either  a  free  acid,  or  salts 
possessing  an  acid  reaction.   [Either  certain  normal  salts,  e.g.  copper 
sulphate,  alum,  etc.,  or  certain  "  acid  "  salts,  as  hydrogen  sodium 
sulphate.     It  must  be  remembered  that  many  of  the  so-called 
"  acid  "  salts,  that  is,  salts  in  which  the  whole  of  the  replaceable 
hydrogen  has  not  been  exchanged  for  metal,  are  not  acid  in  the 
sense  of  exhibiting  an  acid  reaction  towards  litmus  ;   some  are 
neutral,  while  others  are  alkaline  ;  e.g.  hydrogen  sodium  carbonate, 
hydrogen  disodium  phosphate,  etc.]. 

(3)  If  alkaline,  the  liquid  may  contain  either  free  alkali,  or  salts 
having  an  alkaline  reaction. 

B.  When  the  substance  under  analysis  is  a  solid,  it 
should  be  critically  examined,  in  order,  if  possible,  to  gain  any 


Preliminary  Tests  in  a  Systematic  Analysis.     179 

information  from  its  general  physical  properties  which  may  help  to 
identify  it.  If  crystalline,  the  colour,  shape,  etc.,  of  the  crystals 
should  be  noted.  If  powdered,  it  may  be  examined  with  a  pocket- 
lens,  in  order  to  discover  whether  or  not  it  is  homogeneous  ;  i.e. 
whether  it  is  a  single  compound,  or  a  mixture  of  more  than  one. 

The  substance  should  then  be  subjected  to  the  following  general 
tests  * : — 

I.  The  Flame  Reaction. — A  small  quantity  of  the  substance 
is  introduced  into  the  edge  of  the  Bunsen  flame  (first  in  the  cooler 
region  near  the  base  of  the  flame,  and  afterwards  in  the  hotter 
parts  near  the  top  of  the  interior  cone)  upon  a  loop  of  clean  platinum 
wire.     By  a  close  observation  of  the  colour  imparted  to  the  flame, 
compounds   of   sodium,  potassium,    lithium,    barium,  strontium 
calcium,  copper,  and  boron  may  be  indicated  (see  Reactions  of  these 
metals).    Borates  and  certain  thio-cynates  {e.g.  of  mercury)  behave 
in  characteristic  fashion  when  heated. 

II.  The  Borax  Bead.— By  heating  a  minute  particle  of  the 
substance  in  a  borax  bead  in  the  blowpipe  flame,  both  inner  and 
outer  flame,  indications  may  be  obtained  of  the  presence  of  com- 
pounds   of   cobalt,  nickel,    manganese,    iron,    chromium,  copper. 
[Should  this  test  lead  to  the  suspicion  that  either  manganese  or 
chromium  is  present,  it  should  be  followed  up  by  the  fusion  of  a 
portion  of  the   substance  with  sodium  carbonate  and  nitre  upon 
platinum  foil.     See  Manganese  and  chromium  reactions.] 

III.  The  Blowpipe  Reaction  upon  Charcoal,    (a)  When 
heated  alone. — A  considerable    amount   of    information   may   be 
obtained  by  heating  a  little  of  the  substance  by  itself  upon  charcoal. 

(1)  If  it  melts  and  is  absorbed  into  the  charcoal,  it  points  to  the 
substance  consisting  of  salts  of  the  alkalies.  [Chlorates  and  nitrates 
cause  vivid  combustion  of  the  charcoal.] 

(2)  If  a  white  infusible  residue  is  obtained,  the  substance  may 
consist  of  oxides  (or  salts  which  yield  oxides  when  heated)  of  the 
alkaline  earth  metals,  alumina,  zinc  (ZnO  yellow  while  hot,  white 
when  cold),  or  of  silica.     [If  the  residue,  when  placed  upon  a  piece 

*  Every  smallest  detail  of  these  preliminary  tests  should  be  carefully  and 
systematically  noted  down,  whether  the  interpretation  of  the  observed  phe- 
nomena is  obvious  to  the  student  or  not.  Indeed,  it  is  often  just  those  very 
points  which  at  the  time  do  not  convey  any  definite  information  to  him  which 
may  prove  to  be  most  important. 

The  examples  which  are  here  given  of  the  changes  that  may  take  place  when 
substances  are  submitted  to  the  various  general  tests,  must  be  regarded  merely 
as  examples.  It  would  be  obviously  impossible  to  describe  and  tabulate  every 
result  which  might  be  obtained  by  subjecting  the  infinite  variety  of  possible 
mixtures  to  the  tests  described,  and  even  if  it  were  possible,  it  would  be  quite 
undesirable  to  do  so. 


l8o  Qualitative  Analysis. 

of  turmeric  or  litmus  paper  and  moistened  with  water,  shows  an 
alkaline  reaction,  it  will  contain  one  of  the  alkaline  earth  metals. 
Silica  may  be  specially  tested  for  by  the  bead  of  microcosmic  salt.] 

(3)  If  a  coloured  residue  is  left,  it  points  to  one  of  the  metals 
already  indicated  by  the  borax-bead  test. 

(4)  If  reduction  takes  place,  resulting  in  the  formation  of  fumes 
and  an  incrustation  upon  the  charcoal,  without   indication  of  a 
metallic  bead,  it  points  to  the  presence  of  compounds  of  very  volatile 
metals,  as  arsenic  (white  incrustation  accompanied  by  garlic  odour), 
cadmium  (red-brown  incrustation),  zinc  (incrustation  yellow  when 
hot,  and  forming  very  close  to  the  substance).     Such  a  volatile 
compound  as  ammonium  chloride  gives  white  fumes  and  an  in- 
crustation (the  latter  being  formed  at  a  considerable  distance  from 
the  heated  spot).     In  this  case  the  blowpipe  flame  appears  of  a 
yellow-ochre  colour  as  it  impinges  upon  the  ammonium  salt. 

(ft)  When  heated  with  Reducing  Agents. — When  the  substance 
is  mixed  with  sodium  carbonate  and  potassium  cyanide,  and  heated 
upon  charcoal  in  the  reducing  flame,  compounds  of  copper,  silver, 
and  gold  are  reduced  to  the  metallic  state  without  giving  any  in- 
crustation upon  the  charcoal ;  while  compounds  of  antimony, 
bismuth,  tin,  and  lead  give  metallic  beads  accompanied  by 
incrustations  (for  the  characteristics  of  the  incrustations  and  the 
metals,  see  Special  reactions  of  each). 

IV.  The  Action  of  Heat. — The  behaviour  of  a  substance 
when  heated  alone  in  a  dry  narrow  tube  closed  at  one  end  (a  small 
test-tube),  will  generally  afford  important  information  respecting  its 
composition. 

(a)  IF  THE  SUBSTANCE  SIMPLY  MELTS,  AND  SOLIDIFIES  AGAIN 
ON  COOLING,  without  giving  off  any  gases  or  vapours,  a  large  number 
of  compounds  are  obviously  at  once  excluded.  In  this  case  the 
substance  may  consist  of  salts  of  the  alkalies  or  the  alkaline  earth 
metals  ;  a  few  white  salts  of  the  heavy  metals,  e.g.  silver  chloride  ; 
or  a  few  coloured  salts,  e.g.  lead  chromate. 

(£)  IF  WATER  is  GIVEN  OFF,  collecting  in  drops  upon  the  upper 
part  of  the  tube,  it  may  be  due  (i)  to  hygroscopic  *  moisture  (in  this 
case  the  amount  will  probably  be  small),  (2)  to  the  decomposition 
of  metallic  hydroxides,  or  (3)  to  water  of  crystallisation.  [Sub- 
stances containing  much  water  of  crystallisation,  when  heated,  often 
melt  twice—  first  in  the  water  of  crystallisation,  and  as  this  is 

*  That  is  moisture  due  to  the  substance  being  "damp"  —  mechanically 
adhering  moisture.  Almost  all  powders  attract  a  little  moisture  in  this  way 
from  the  atmosphere. 


Preliminary  Tests  in  a  Systematic  Analysis.     181 

expelled  they  resolidify,  but  on  the  application  of  a  stronger  heat 
they  once  more  undergo  fusion.] 

The  water  should  be  carefully  tested  by  introducing  a  small 
strip  of  litmus  paper. 

Alkalinity  would  indicate  ammonium  compounds  (e.g. 
(NH4)2MgPO4,  HNa(NH4)PO4,  etc.). 

Acidity  might  be  due  to  the  decomposition  of  certain  acid 
salts,  either  alone  or  by  interaction  with  other  salts  (e.g.  hydrogen 
potassium  sulphate,  when  heated,  is  converted  into  the  normal 
salt  and  sulphuric  acid,  and  if  present  along  with  a  salt  containing 
water  of  crystallisation,  the  water  would  become  acid  ;  or  if  mixed 
with  sodium  chloride  or  nitrate,  hydrochloric  acid  or  nitric  acid 
would  be  evolved  by  double  decomposition). 

(c)  IF  THE  SUBSTANCE  CHANGES  COLOUR- 
CO  From  white  to  yellow,  indicates  oxide  of  zinc,  tin,  bismuth  ; 
(2)  From  yellow  to  brown  (fusing  at  a  red  heat),  oxide  of  lead. 
Most  coloured  oxides  become  much  darker  when  heated  (e.g. 

HgO,  Fe203)- 

(d)  IF  GASES  OR  VAPOURS  ARE  EVOLVED,  they  must  be  iden- 
tified by  special  tests. 

(1)  Oxygen  and  nitrous  oxide  (re-ignition  of  glowing  splint  of 
wood).     The  former  from  chlorates,  iodates,  bromates,  nitrates,  or 
peroxides  ;  the  latter  from  ammonium  nitrate,  or  salts  which  by 
double  decomposition  give  ammonium  nitrate. 

(2)  Hydrogen  (character  of  flame),  from  formates,  and  from 
oxalates     in    presence    of    caustic    alkalies    or    alkaline    earths 
(K2C2O4  +  2KHO  =  2K2CO3  +  H2). 

(3)  Nitrogen  (extinguishes  flame,  and  gives  no  reaction  with 
lime-water),  from  ammonium  nitrite  or  chromate,  or  from  mixtures 
which  yield  these  salts  by  double  decomposition. 

(4)  Chlorine,  bromine,  and  iodine  (recognised  by  colour,  smell, 
etc.),  evolved  from  certain  halogen  compounds  when  heated  alone, 
or  in  admixture  with  acid  salts  and  peroxides  (e.g.  a  mixture  con- 
taining such  compounds  as  Nad,  HKSO4,  and  MnO2,  when  heated, 
evolves  chlorine). 

(5)  Sulphur  vapour  (sublimate  of  yellow  drops),  from  metallic 
persulphides. 

(6)  Carbon  dioxide  (action  on  lime-water),  from  most  carbonates 
other  than  the  normal  salts  of  the  alkalies  (also,  mixed  with  carbon 
monoxide,  from  certain  formates  and  oxalates). 

(7)  Sulphur  dioxide  (characteristic  odour),  from   certain   sul- 
phites (e.g.  CaSO3  =  CaO  +  SO2),  certain  sulphates  (e.g.  2FeSO4 


1 82  Qualitative  Analysis. 

-  Fe2O3  +  S03  +  SO., ;  CuSO4  =  CuO  +  SO2  +  O).  It  may  also 
be  due  to  the  oxidation  of  sulphur  evolved  from  sulphides  or  other 
sulphur  compounds,  as  thiocyanates. 

(8)  Cyanogen   (formation  of  thiocyanate  of  ammonia),  from 
certain  metallic  cyanides. 

(9)  Ammonia  (odour,  and  action  on  test-paper),  from  many 
ammonium  salts ;  e.g.  sulphate,  phosphate,  carbonate,  etc.  (also 
from  the  decomposition  of  cyanates). 

(10)  Sulphuretted    hydrogen    (odour,    and    action    on    paper 
moistened  with  lead  acetate),  from  the  decomposition  of  hydrated 
metallic  sulphides,  of  sulphites  and  thiosulphates,  in  the  presence 
of  compounds  which  yield  water  or  acid. 

(n)  Phosphoretted  hydrogen  (odour,  and  character  of  flame), 
from  hypophosphites. 

(12)  Nitrogen  peroxide  and  oxygen,  from  the  decomposition 
of  nitrates  of  heavy  metals,  or  from  mixtures  which  by  double 
decomposition  yield  such  a  compound  {eg.  PbSO4  and  KNO3\ 

0?)  IF  A  SUBLIMATE  is  PRODUCED,  it  indicates  such  volatile 
solids  as  the  following  : — 

(1)  Giving  a  white  sublimate— ammonium  haloid  salts,  arsenious 
oxide,  mercuric  chloride,  mercurous  chloride  (yellowish  while  hot, 
white  on  cooling). 

(2)  Giving  a  coloured   sublimate — mercuric   iodide   (red  and 
yellow),  arsenious  sulphide  (yellow). 

(3)  Giving  a  black  metallic-looking  sublimate — iodine    (violet 
vapours),  mercuric  sulphide  (red  streaks  if  rubbed  with  a  glass  rod), 
metallic  mercury  (runs  into  liquid  globules  when  rubbed  with  a 
glass  rod),  metallic  arsenic  (garlic  smell).     [If  this  test  leaves  it 
doubtful  whether  arsenical  or  mercurial  compounds  are  present, 
a  little  of  the  substance  should  be  mixed  with  sodium  carbonate 
and  heated  in  a  dry  tube.     If  mercurial,  the  sublimate  consists  of 
minute  globules.] 

Ammoniacal  compounds  may  be  at  once  tested  for  by  heating 
a  portion  of  the  substance  with  a  little  sodium  hydroxide,  when 
ammonia  is  evolved. 

(/)  IF  NO  CHANGE  TAKES  PLACE,  it  will  be  evident  from  the 
foregoing  that  the  number  of  substances  which  can  possibly  be 
present  is  extremely  limited  ;  and  if  in  addition  the  compound  is 
white,  the  range  of  probable  substances  is  still  further  narrowed 
down  to  the  oxides  of  a  few  of  the  metals,  such  as  the  alkaline 
earths,  alumina,  silica,  etc. 

To  obtain  a  Solution  of  the  Solid  Substance,    (i)  In 


Solution  of  a  Solid  Substance  183 

Water. — A  small  quantity  of  the  substance,  in  a  finely  powdered 
condition,  is  first  treated  with  cold  water  in  a  large  test-tube. 
[This  may  result  in  the  evolution  of  gas,  owing  to  the  chemical 
action  of  water  upon  such  compounds  as  certain  metallic  phosphides 
(e.g.  calcium  phosphide)  or  carbides  (calcium  carbide),  which 
evolve  phosphoretted  hydrogen  and  acetylene  respectively;  or 
carbon  dioxide  might  be  expelled,  owing  to  the  presence  of  soluble 
carbonates  and  acid  salts,  which  would  interact  upon  each  other.] 
The  mixture  is  then  boiled  for  a  few  moments. 

If  the  substance  does  not  dissolve,  it  is  either  wholly  or  partly 
insoluble  in  water.  To  ascertain  whether  any  of  it  has  dissolved, 
the  mixture  should  be  allowed  to  settle,  and  a  few  drops  of  the 
clear  liquid  evaporated  to  dryness  upon  a  watch-glass  ;  or  upon  a 
piece  of  platinum  foil,  cautiously  heated.  If  a  residue  is  obtained 
on  evaporation  (thus  showing  that  partial  solution  in  water  has 
taken  place),  the  aqueous  liquid  is  decanted  off  (or  filtered  if  the 
insoluble  part  does  not  readily  settle),  and  the  insoluble  portion 
again  boiled  with  water  and  filtered. 

(2)  In  Acids.— The  portion  insoluble  in  water  is  then  treated 
with  dilute  hydrochloric  acid  and  boiled.  If  it  does  not  entirely 
dissolve,  the  dilute  acid  is  decanted  off  into  another  test-tube,  and 
strong  hydrochloric  acid  substituted,*  the  mixture  being  submitted 
to  prolonged  boiling,  if  necessary.  If  the  substance  wholly  dissolves, 
the  two  acid  solutions  may  be  mixed  together. 

If,  on  adding  a  few  drops  of  the  aqueous  extract  to  a  small 
portion  of  the  acid  solution,  no  precipitation  takes  place,  the  main 
portions  of  the  two  solutions  may  be  mixed  together  f  and  examined 
for  metals  and  acids  (except  hydrochloric  acid)  by  the  methods 
already  described.  On  the  other  hand,  if  the  two  extracts  contain 
compounds  which  will  interact  with  the  formation  of  an  insoluble 
precipitate,  they  must  be  examined  separately. 

If  the  substance  is  not  entirely  dissolved  by  strong  hydrochloric 
acid,  it  may  consist  of  one  of  the  few  compounds  which  are  only 
dissolved  by  aqua  regia.\  The  residue  is  therefore  treated  with  a 
small  quantity  of  mixed  nitric  and  hydrochloric  acids,  and  boiled. 
The  solution  should  be  evaporated  down  in  a  porcelain  dish  until 

*  By  the  behaviour  of  the  substance  under  this  treatment,  the  presence  of 
carbonates,  sulphites,  and  peroxides  will  be  indicated.  Sulphur  and  silicic 
acid  may  be  precipitated. 

t  See  remarks  on  p.  187  with  reference  to  this  point. 

j  Such  as  sulphides  of  nickel,  cobalt,  mercury.  The  presence  of  these  should 
have  been  indicated  by  the  preliminary  tests.  The  use  of  aqua  regia  should 
only  be  resorted  to  when  absolutely  necessary. 


184  Qualitative  Analysis. 

only  a  small  bulk  remains,  in  order  to  expel  as  much  acid  as 
possible,  and  then  diluted  with  water ;  this  solution  may  then  be 
mixed  with  the  hydrochloric  acid  extract. 

(3)  Treatment  of  the  Residue  insoluble  in  Water  or 
in  Acids. — The  number  of  compounds  insoluble  in  water  and 
acids  is  very  limited.  Of  commonly  occurring  substances,  the 
following  may  be  present — 

(a)  Sulphates  of  barium,  calcium,  strontium,  and  lead  (the  three 
last  named  being  sufficiently  soluble  to  be  detected  in  the  aqueous 
extract). 

(^)  Silica,  and  many  natural  silicates. 

(c)  Fluor  spar,  and  other  natural  fluorides. 

(d)  Silver   chloride  (such  insoluble   silver   compounds   as   the 
bromide,  iodide,  cyanide,  and  ferrocyanide  are  converted  into  the 
chloride  by  boiling  with  aqua  regia). 

(e)  Oxides    of   aluminium   and   chromium   (which   have   been 
strongly  heated).     Native  stannic  oxide  (tinstone.} 

(y)  A  few  arsenates  and  phosphates.     Carbon  and  sulphur. 

These  insoluble  substances  (except  carbon  and  sulphur,  which 
are  readily  recognised  and  require  no  further  treatment)  are  con- 
verted into  soluble  compounds  by  fusion  with  alkali  carbonates.* 
The  insoluble  residue  (in  the  absence  of  lead  and  silver)  is  dried, 
and  mixed  with  about  four  times  its  weight  of  fusion  mixture  and  a 
little  potassium  nitrate,  and  fused  in  a  platinum  crucible  until  all 
effervescence  is  at  an  end.  The  crucible  is  then  allowed  to  cool. 
[If  the  crucible  be  quickly  chilled  by  standing  it  upon  a  piece  of 
cold  iron,  e.g.  the  foot  of  a  retort  stand,  or  by  dipping  it  on  to  the 
surface  of  cold  water  in  a  basin,  the  solidified  contents  of  the  crucible 
at  once  crack  away  from  the  metal,  and  can  be  removed  as  a  solid 
cake,  which  can  be  broken  up  in  a  mortar.] 

The  fused  mass  is  then  boiled  with  water  until  nothing  further 
dissolves,  and  filtered. 

The  aqueous  solution  is   examined  for  such   acids  as   in  the 

*  If  lead  sulphate  or  silver  chloride  be  present,  they  must  either  be  re- 
moved, or  the  fusion  must  be  conducted  in  a  crucible  of  porcelain  instead  of 
platinum  (in  which  case  a  certain  amount  of  alumina  from  the  crucible  will 
pass  into  combination  with  the  alkali  carbonate).  Lead  sulphate  may  be  re- 
moved by  boiling  the  insoluble  residue  in  a  strong  solution  of  ammonium 
acetate  (the  solution  being  tested  for  lead  and  for  sulphuric  acid).  Silver 
chloride  may  be  dissolved  out  by  a  solution  of  potassium  cyanide  ;  or  it  can  be 
reduced  to  metallic  silver  by  means  of  zinc  (see  Silver,  reaction,  p.  114),  and 
after  washing  free  from  zinc  chloride,  the  reduced  silver  can  be  dissolved  in 
nitric  acid.  If  sulphur  be  present  in  the  insoluble  residue,  it  may  be  expalled 
by  heating  the  substance  in  a  porcelain  crucible. 


Treatment  of  Insoluble  Substances.  185 

nature  of  the  case  could  be  present — namely,  sulphuric,  silicic, 
hydrofluoric,  chromic,  arsenic,  phosphoric,  hydrochloric.* 

The  residue,  after  being  thoroughly  washed  with  hot  water, 
until  the  wash-water  is  no  longer  alkaline,  is  dissolved  in  hydro- 
chloric acid,  and  the  solution  examined  for  such  metals  as  can  be 
present. 

In  a  few  special  cases,  fusion  with  hydrogen  potassium  sulphate 
is  resorted  to  instead  of  the  fusion  mixture  above  mentioned.  This 
method  is  applicable,  for  example,  in  the  case  of  minerals  contain- 
ing titanium,  such  as  titaniferous  iron  ores  (see  pp.  73,  343). 

*  The  fusion-mixture  must  obviously  be  free  from  sulphates,  chlorides,  etc.. 
or  tests  for  these  acids  are  of  no  value. 


GENERAL  TABLE  FOR  THE  SEPARATION  OF  METALS 

To  the  solution  of  the  substance  under  analysis  (see  pp.  182,  et  seq.}  add  a  few 

effecting  the  solution  of  the  substance,  this  first  step  in  the  general  separation  is 

until  the  precipitation  is  complete.     Warm  gently  (see 


The    Precipitate 

may  consist  of 
AgCl 
Hg2Cl2 
PbCl2 

Examine  accord- 
ing to  the  table 
on  p.  115 

Group  I. 


The  Filtrate  is  gently  warmed,  and  a  stream  of  sulphuretted 
tation  is  complete. 

[NOTE. — To  ensure  complete  precipitation,  it  is  necessary  that  the  solution 
when  this  element  is  known  to  be  present,  many  chemists  boil  the 
(any  sulphate  of  lead  or  barium  which  ma)'  thereby  be  precipitated  is 
before  sulphuretted  hydrogen  is  passed  into  it.] 

A  small  portion  of  the  mixture  should  be  filtered,  and  the 
hydrogen  passed  into  it.  Should  any  further  precipitation 
treated  to  more  gas. 


The  Precipitate  may  consist 
of— 

(1)  PbS  ;  HgS  ;  Bi2S3 ;  CuS  ; 

CdS. 

(2)  Sb2S3 ;  A?2S3 ;  SnS  ;  SnS2 

Wash  thoroughly,  and  then 
transfer  it  to  a  small  beaker 
and  warm  gently  with 
yellow  ammonium  sulphide, 
p.  107.  [If  mercury  is 
known  to  be  absent,  caustic 
soda  may  be  used  instead 
of  the  sulphide.] 


The    Besidue     The    Filtrate 

may  contain  ;  may  contain 
the  sulphides  the  thio  salts 
of  Division  I.  of  As,  Sb, 


to 


_.    . 

Examine 
cording 
the   table  on 
p.  88. 


Group  II., 
Div.    i 


Examine  ac- 
cording to 
the  table  on 
p.  108. 


Group  II., 
Div.  2. 


The  Filtrate.     Boil  until  sul- 

boil  again  for  a  few  moments 

hydrogen),  and  then  evapo- 

it  has   become  diluted.     If 

ration  must  be  carried  down 

the   silica   insoluble.      The 

Test  a  small  portion  of  the 

To  the  main  portion  of  the 

until   precipitation   is  com- 


The  Precipitate. 

I.  In  the  absence  of  phos- 
phoric acid,  may  consist 
of— 

A12(HO)6;  Cr2(HO)6; 
Fe2(HO)6. 

Examine  by  table  on  p.  49. 


Group  IIlA. 


II.  In  the  presence  of  phos- 
phoric acid,  besides  the 
above  hydroxides,  the  pre- 
cipitate may  contain  the 
phosphates  of  any  or  all  of 
the  metals  of  Groups  III. 
and  IV.,  and  of  magnesium. 

Examine  by  special  table  on 
following  page. 


INTO  GROUPS  (see  also  pp.  16,  17,  and  116,  117). 

drops  of  dilute  HC1.  [Obviously  in  cases  when  HC1  has  already  been  employed  in 
omitted.]  If  any  precipitate  is  produced,  continue  adding  the  reagent  gradually 
footnote,  p.  107),  then  thoroughly  cool  again,  and  filter. 

hydrogen  slowly  bubbled  through  the  liquid,  with  frequent  stirring,  until  precipi- 


be  not  too  strongly  acid — see  Cd,  p.  83  ;  Pb,  p.  80 ;  and  As,  p.  93.     JLU   me  \M 
solution  with  sulphurous  acid  to  ensure  the  rapid  reduction  of  the. arsenic  to  the  a 


vith  sulp 
filtered  off  and  treated  separately). 


In  the  case  of  arsenic, 
rsenious  state 
The  solution  must  be  boiled  to  expel  all  the  sulphur  dioxide 


filtrate  diluted  with  two  or  three  times  its  volume  of  water,  and  more  sulphuretted 
result,  the  main  portion  must  be  similarly  diluted,  without   being  filtered,  and 


phuretted  hydrogen  is  entirely  expelled.  Add  two  or  three  drops  of  HNO3,  and 
(to  oxidize  any  iron  or  chromium  which  may  have  been  reduced  by  the  sulphuretted 
rate  the  liquid  to  about  half  its  volume,  or  less,  according  to  the  extent  to  which 
silica  or  organic  compounds  have  been  detected  by  preliminary  tests,  the  evapo- 
to  dryness,  and  the  residue  gently  heated  in  order  to  destroy  the  latter,  and  render 
residue  is  then  extracted  with  HC1  and  water  (see  p.  117). 
solution  for  phosphoric  acid  (see  pp.  63  and  117). 

solution  add  a  considerable  quantity  of  NH4C1,  and  heat  to  boiling.    Add  NH4HO 
plete  (see  p.  48),  boil  for  a  moment,  and  filter  while  hot. 


The  Filtrate.     Pass  sulphuretted  hydrogen  (or  add  ammonium   sulphide)  until 
precipitation  is  complete.     Gently  warm  the  liquid  (see  Ni,  p.  57),  and  filter. 


The    Precipitate  <  The  Filtrate.     Boil  to  expel  H2S.     If  ammonium  sulphide 
may  consist  of —        has  been  employed,  add  a  little  HC1  before  boiling.     If 
necessary,   concentrate  the  solution  by  evaporation.     Add 
NH4HO  until  alkaline,  and  (NH4)2CO3  until  precipitation 
is  complete.     Warm  the  liquid,  but  do  not  boil  (see  p.  34). 


MnS  ;  ZnS  ;  NiS 
CoS. 


Examine  by  table 
on  p.  62. 


Group  IIlB. 


The    Precipitate 
may  consist  of — 

BaCO3 ;  SrCO3 ; 
CaCO3. 

Examine  by  table 
on  p.  35. 

Group  IV. 


The  Filtrate, 

Examine  for  Mg,  K,  Na,  as  directed 
on  p.  25. 

Group  V. 


NOTE  i. — In  testing  for  Mg  at  this  point,  it  will 
obviously  be  unnecessary  to  add  more  NH4C1 
and  NH4HO. 

NOTE  2. — Any  failure  to  effect  complete  group 
separations  will  usually  result  in  the  precipita- 
tion of  some  metallic  phosphate  af  this  stage, 
other  than  magnesium  phosphate. 


TABLE  FOR  GROUP  III.,  PHOSPHORIC 


The  precipitate  produced  by  NH4HO  in  presence  of  NH4C1  may  consist 
of  Groups  III.  and  IV.,  and  of  Mg.      Dissolve  in  a  little  dilute  HC1, 

and  add  sodium  acetate  and  acetic  acid  * 


The  Precipitate  may  consist  of  phos- 
phates of  Fe,  Al,  Cr  (along  with 
basic  acetate  of  iron). 

Examine  by  table  on  p.  49 
(see  also  p.  69). 


The  Filtrate.  To  a  small  portion 
If  a  precipitate  is  produced, 
drop,  until  the  whole  of  the 
the  precipitation  is  seen  by  the 
mixture,!  and  filter. 

If  the  preliminary  test  with  ferric 
contains  no  more  phosphoric 


The  Precipitate  consists  of  ferric 
phosphate  and  basic  acetate 


Throw  away. 


*  If  the  addition  of  sodium  acetate  gives  no  precipitate,  it  follows 
analysis  (see  footnotes,  pp.  68,  69). 

If,  after  the  addition  of  the  acetate  reagent,  and  before  the  mixture 
allowing  the  precipitate  to  settle  somewhat),  this  means  that  ferric  acetate 
solution)  ;  and  if  this  is  formed,  it  follows  that  there  is  present  in  the 
present ;  hence  all  the  phosphoric  acid  will  have  passed  into  the  precipitate 

f  Boiling  the  mixture  at  this  point  insures  the  conversion  of  the  soluble 
solution  should  be  colourless,  or  entirely  free  from  the  red  colour. 


ACID   BEING   PRESENT   (see  also  p.  69). 


of  hydroxides  of  Al,  Fe,  Cr,  as  well  as  phosphates  of  any  of  the  metals 
avoiding  excess.  Nearly  neutralize  with  Na2CO3  (see  footnote,  p.  68), 
(reagent).  Boil  the  mixture,!  and  filter. 

of  the  liquid,  which  should  be  colourless,  add  a  drop  of  ferric  chloride, 
then  ferric  chloride  is  added  to  the  main  portion  of  the  filtrate,  drop  by 
phosphoric  acid  is  thrown  down  as  ferric  phosphate.  The  completion  of 
liquid  becoming  red,  by  the  formation  of  ferric  acetate.  Boil  the 

chloride  gave  no  precipitate,  but  only  a  red  colouration,  then  the  solution 
acid,  and  is  at  once  treated  as  in  the  next  step. 


The    Solution.      Add    NH4C1,    heat    to   boiling,    and    add    NH4HO. 
Filter. 


The    Precipitate    may 

consist    of    A12(HO)6 
and  Cr2(HO)6. 


The  Filtrate. 
Examine  for  Groups  IIlB.  and  IV.,  and  for 


Magnesium  in  the  usual  way. 

Examine  in  usual  way. 

I 


that  no  iron,  aluminium,  or  chromium  are  present  in  the  substance  under 

has   been  boiled,   the  liquid  itself  appears   red  (which   is   easily  seen  by 
is  being  produced  (which    is   soluble   in   sodium   acetate   forming   a   red 
mixture   more  than  enough  iron   to   unite    with    all   the  phosphoric   acid 
along  with  the  iron, 
ferric  acetate  into  the  insoluble  basic   acetate,  so  that  when    filtered  the 


CHAPTER  XVIT. 
THE  RESULTS  OF  A  QUALITATIVE  ANALYSIS. 

AT  the  end  of  an  analysis,  the  student  will  be  in  possession  of  the 
knowledge  that  the  substance  he  has  examined  contained  com- 
pounds of  certain  metals  or  positive  radicals,  and  certain  negative 
or  acid  radicals.  If  only  one  metal  and  one  acid  have  been  found, 
the  substance  must,  of  course,  be  the  compound  of  these  two  ;  e.g.  if 
magnesium  and  sulphuric  acid  have  been  detected,  the  compound 
must  have  been  magnesium  sulphate.  If  a  metal  has  been  found, 
but  no  acid  radicals,  the  substance  will  in  all  probability  be  an 
oxide  of  that  metal.  In  deciding  this  point,  however,  the  student 
must  consider  the  general  properties  of  the  substance,  for  if  he 
finds,  for  example,  the  metal  magnesium  in  a  substance  soluble  in 
water,  the  compound  obviously  cannot  be  the  oxide.  He  must 
therefore  conclude  that  he  has  overlooked  the  acid  radical. 

When  more  than  two  metals  and  acid  radicals  are  found  in  the 
substance,  it  becomes  a  much  more  difficult  matter  to  discover 
which  acid  is  in  combination  with  which  metal.  Indeed,  in  many 
cases  it  is  not  possible  to  decide  this  point  by  a  qualitative  analysis 
alone.  For  example,  suppose  a  mixture  is  found  on  analysis  to 
contain  the  metals  cadmium  and  sodium,  and  the  acid  radicals  of 
sulphuric  and  nitric  acids.  Does  the  mixture  consist  of  cadmium 
nitrate  and  sodium  sulphate,  or  of  cadmium  sulphate  and  sodium 
nitrate?  All  the  four  salts  are  white,  and  none  of  them  is  so 
markedly  different  from  the  others  in  its  solubility  in  water  as  to 
throw  any  light  on  the  question  ;  so  that,  if  the  two  salts  are  present 
in  fairly  equal  quantities,  it  is  impossible  to  decide  which  pair  of 
salts  is  present. 

There  are  a  number  of  points,  however,  which  a  student  who 
cultivates  his  powers  of  observation  will  be  able  to  note,  which 
will  often  give  a  clue,  if  they  do  not  afford  proof,  as  to  the  manner 
in  which  the  metals  and  acid  radicals  are  united  in  a  mixture. 

For  instance,  even  in  such  a  mixture  as  the  above,  by  carefully 


The  Results  of  a  Qualitative  Analysis.          187 

noting  the  relative  quantities  of  various  precipitates  obtained  during 
the  course  of  analysis,  some  light  might  be  thrown  on  this  point. 
Suppose  it  was  noted  that  the  amount  of  cadmium  sulphide  thrown 
down  seemed  very  small,  and  that  the  test  for  sulphuric  acid  re- 
sulted in  a  most  copious  precipitation  of  barium  sulphate,  it  would 
be  fair  to  conclude  that  the  cadmium  was  not  present  as  sulphate. 

Again,  suppose  barium  and  ammonium  are  the  "  metals  "  found 
along  with  the  acid  radicals  of  hydrochloric  and  nitric  acids.  The 
four  possible  salts  are  in  this  case  also  all  of  them  white  and  easily 
soluble.  If  the  salts  are  in  fairly  equal  quantities,  is  it  possible  to 
decide  how  the  acids  and  metals  are  distributed  ?  If  the  dry  re- 
actions had  been  carefully  observed,  they  would  enable  the  student 
to  decide  the  point.  A  mixture  consisting  of  BaCl2  and  NH4NO3 
when  heated  would  melt,  and  give  off  water  and  nitrous  oxide 
(decomposition  of  the  ammonium  nitrate),  whereas  if  Ba(NO3)2  and 
NH4C1  were  present,  some  of  the  ammonium  chloride  would  at 
once  sublime  up  the  tube  ;  and  later  on,  brown  fumes  would  appear, 
due  to  the  decomposition  of  barium  nitrate. 

Very  often  two  colourless  and  soluble  salts  contained  in  a 
mixture,  undergo  double  decomposition  with  formation  of  a  pre- 
cipitate, as  soon  as  water  is  added  to  the  mixture  with  a  view  to  its 
solution.  When  the  precipitate  is  coloured  (as  would  be  the  case, 
for  example,  if  the  mixture  consisted  of  silver  nitrate  and  sodium 
phosphate),  the  fact  that  it  has  been  formed  during  the  process  of 
solution,  and  was  not  originally  present,  will  be  evident.  On  the 
other  hand,  when  the  product  of  interaction  is  also  white,  it  is  not 
always  possible  to  decide  this  point.  If  such  a  mixture  as  silver 
nitrate  and  barium  chloride  be  treated  with  water,  the  two  soluble 
5alts  interact,  with  the  precipitation  of  silver  chloride.  In  this  case 
a  close  inspection  of  the  appearance  of  the  mixture  (the  milkiness 
imparted  to  the  water)  would  indicate  that  the  substances  were 
interacting.  On  the  other  hand,  with  a  mixture  such  as  barium 
chloride  and  magnesium  sulphate,  it  would  be  almost  impossible  to 
say  whether  the  original  mixture  contained  barium  sulphate  or  not. 

When  the  substance  under  analysis  is  partly  soluble  in  water 
and  partly  soluble  in  acids,  it  will  be  obvious  that  much  informa- 
tion as  to  the  distribution  of  the  metals  and  acids  will  be  gained,  if 
each  solution  is  separately  examined.  Thus,  in  a  mixture  consist- 
ing of  lead  carbonate  and  barium  nitrate,  water  will  remove  the 
latter  salt,  and  it  becomes  easy  to  settle  that  the  lead,  and  not  the 
barium,  is  in  combination  with  the  carbonic  acid. 

These  are  simple  illustrations  of  the  kind  of  evidence  which  has 


1 88  Qualitative  Analysis. 

to  be  made  use  of  in  ordinary  qualitative  analysis,  when  endeavour- 
ing to  settle  the  question  of  the  distribution  of  the  metals  and  acid 
radicals.  It  will  be  obvious  that  the  success  with  which  the  analyst 
will  meet  in  this  part  of  his  work  will  largely  depend  upon  the  power 
which  he  possesses  of  seeing  (i.e.  of  observing)  all  the  various 
indications  which  are  given  him  during  the  course  of  the  analysis. 
The  chief  aim  of  the  analytical  student,  as  it  should  be  the 
first  object  of  his  instructor,  must  always  be  the  cultivation  and 
development  of  these  powers  of  exact  observation.  Without  this, 
anylysis  sinks  into  a  mere  dull  mechanical  routine,  utterly  devoid 
of  all  interest  or  educational  value- 


BOOK   II. 

QUANTITATIVE   ANALYSIS. 

PART  I. 

GRAVIMETRIC  METHODS. 

SECTION  I. 
PRELIMINARY  MANIPULATIONS. 

i.  Weighing. — In  gravimetric  methods  of  analysis,  as  the  name 
indicates,  the  quantities  which  it  is  the  object  of  the  analysis  to 
determine,  are  estimated  by  the  process  of  weighing.  Not  only  is 
the  amount  of  the  material  employed  for  analysis  a  'weighed  quantity 
(which  may  also  be  the  case  in  volumetric  methods),  but  the  final 
products  of  the  various  processes  are  obtained  in  a  form  in  which 
they  also  can  be  weighed.  For  instance,  in  making  a  quantitative 
analysis  of,  say,  a  silver  coin,  a  weighed  quantity  of  the  metal  will 
be  operated  upon,  and  the  silver  and  copper  will  finally  be  obtained 
in  the  form  of  silver  chloride  and  copper  oxide.  The  exact  quanti- 
ties of  these  products  will  then  be  ascertained  by  the  process  of 
weighing,  and  from  the  weights  so  obtained  the  proportions  of  the 
two  metals  are  deduced  by  calculation. 

In  order  to  obtain  the  accurate  weight  of  a  substance,  it  is 
necessary  to  employ  a  delicate  balance,  and  to  perform  the  opera- 
tions of  weighing  with  careful  manipulation. 

Fig.  15  represents  a  modern  form  of  chemical  balance.  It  is 
contained  in  a  glass  case  with  counterpoised  sliding  sashes,  to 
exclude  from  it  as  much  as  possible,  dust,  damp,  and  laboratory 
fumes  ;  *  and  also  to  prevent  disturbance  due  to  air-draughts  during 
the  process  of  weighing. 

*  Unless  circumstances  render  it  unavoidable,  the  balance  should  never  be 
kept  in  the  laboratory.  In  all  properly  equipped  schools  a  balance-room  is 
provided,  the  instruments  being  placed  where  they  will  be  subject  to  the  least 
possible  vibration,  and  never  being  moved. 

O 


i  go     Quantitative  Analysis  :   Gravimetric  Methods. 

The  case  is  mounted  upon  levelling-screws,  by  means  of  which 
the  instrument  can  be  placed  in  a  perfectly  horizontal  position, 
which  is  indicated  by  the  spirit-levels  attached  to  the  base  of  the 
pillar. 

When  the  balance  is  at  rest,  the  beam  and  the  scale-pans  are 
held  by  the  mechanical  supports,  which  can  be  operated  upon  by 
means  of  the  handle  or  "  milled-head  "  placed  outside  and  in  front 
of  the  case. 


FIG.  15. 


By  a  slight  turn  of  the  milled-head,  these  supports  are  simul- 
taneously lowered,  and  the  beam  and  pans  are  free  to  swing  upon 
their  "knife-edges." 

The  extent  to  which  the  beam  swings  is  indicated  by  a  pointer 
and  scale  ;  and  if  the  balance  is  in  proper  order  and  adjustment,  it 
will  be  seen  that  as  the  beam  moves,  the  pointer  travels  practically 
an  equal  number  of  graduations  in  each  direction. 

*  If  the  pointer  swings  further  in  one  direction  than  in  the  other,  either  the 
balance  is  not  level,  or  is  out  of  adjustment.  The  devices  for  adjusting  the 
instrument  differ  with  different  balances.  In  the  example  shown  in  the  figure, 
it  is  effected  by  means  of  the  two  little  "  bobs  "  which  travel  upon  two  screws 
attached  to  the  ends  of  the  graduated  scale  of  the  beam. 


Preliminary  Manipulations. 


191 


When  not  in  use,  the  balance  must  never  be  left  in  this  condition, 
but  the  supports  must  always  be  raised. 

It  is  desirable  to  keep  inside  the  balance  case  a  glass  jar  about 
half  filled  with  dry  calcium  chloride,*  or  fragments  of  quicklime, 
with  a  view  to  render  the  atmosphere  of  the  case  as  dry  as  possible. 
This  is  especially  necessary  when  any  of  the  vital  working  parts  of 
the  instrument  are  made  of  steel  instead  of  agate. 

A  set  of  gram  weights,  suitable  for  chemical  analysis,  is  seen  in 
Fig.  16.  The  weights  from  I  gram  to  50  grams  (usually  arranged 
to  make  up  100  grams  in  all)  are  of  brass,  while  the  fractions  of 
the  gram  are  flat  pieces  of  platinium  or  aluminium  foil.  In 
practice  it  is  usual  to  employ 
only  the  "  deci "  and  "  centi " 
grams  of  these  small  weights 
(i.e.  the  first  and  second  places 
of  decimals),  and  for  the  milli- 
grams and  fractions  of  a  milli- 
gram (the  third  and  fourth 
decimals)  to  use  a  sliding 
weight  upon  the  beam  of  the 
balance.  For  this  purpose 
the  beam  is  graduated  (in  the 
figure,  into  tenths  and  hun- 
dredths),  and  the  sliding 
weight,  or  rider  (consisting  of 
a  piece  of  gilded  brass  wire, 
bent  so  that  it  can  be  placed  astride  the  beam),  is  moved  to  any 
desired  position  on  the  beam  by  means  of  the  carrier  which  passes 
through  the  side  of  the  balance  case.  The  weight  of  the  rider  is 
one  centigram,  and  when  it  is  placed  on  the  tenth  main  gradua- 
tion, its  effect  is  the  same  as  if  it  were  placed  in  the  scale,  i.e.  10 
milligrams  or  I  centigram.  The  divisions  of  the  beam,  therefore, 
represent  milligrams  and  sub-divisions  of  a  milligram. 

The  weights  must  on  no  account  be  touched  with  the  hand,  but 
should  be  manipulated  entirely  by  means  of  the  forceps  supplied 
with  each  set. 

Great  care  should  be  taken  in  touching  these  parts,  and  it  may  be  accepted 
as  a  general  rule  that  the  young  student  should  not  himself  attempt  the  adjust- 
ment of  the  balance. 

*  The  vessel  should  not  be  filled,  especially  when  calcium  chloride  is  used  ; 
as  the  salt  becomes  liquid  by  the  absorption  of  moisture,  and  is  likely  to  overflow 
the  vessel.  This  is  liable  to  happen  when  the  balance  is  left  unused  for  any- 
lengthy  period,  as,  for  example,  during  a  vacation. 


FIG.  16. 


192     Quantitative  Analysis :    Gravimetric  Methods. 

It  is  essential  that  the  weights  all  bear  the  exact  relation  to  each 
other  which  is  professed  by  the  values  indicated  upon  them ;  that 
is  to  say,  the  i-gram  must  be  exactly  half  the  2-gram  weight,  and 
this  must  be  exactly  a  fifth  of  the  10  gram,  and  so  on.*  To  ascer- 
tain that  this  is  actually  the  case,  they  should  be  tested  against  each 
other — first  by  seeing  that  the  several  i-gram  weights  are  equal  to 
each  other  ;  then  that  two  of  them  exactly  counterpoise  the  2 -gram  ; 
then  that  the  2-gram  and  the  three  i -grams  together  equal  the 
weight  of  the  5 -gram,  and  so  on.f  When  using  the  balance,  the 
object  to  be  weighed  should  be  placed  upon  the  left  scale-pan, 
as  it  is  more  convenient  for  manipulation  to  use  the  right-hand 
pin  for  the  weights  (a  left-handed  operator  will  reverse  this 
order). 

Whenever  anything  (either  objects  to  be  weighed,  or  the  weights) 
is  put  upon  the  scale  or  removed  from  it,  the  beam  must  be  at  rest 
on  its  supports. 

The  actual  operation  is  conducted  as  follows  :  The  article  to  be 
weighed,  say  a  porcelain  crucible  with  its  lid,  is  placed  upon  the 
left  scale-pan,  the  balance  being  at  rest.  A  weight,  which  by  a 
guess  is  judged  to  be  rather  greater  than  that  of  the  object,  is 
deposited  by  means  of  the  forceps  upon  the  opposite  scale,  and 
the  beam  gently  liberated  by  turning  the  milled-head  with  the 
left  hand.  Suppose  the  2o-gram  weight  has  been  selected,  and 
found  too  heavy  ;  it  is  then  removed,  returned  to  its  place  in  the 
box,  and  replaced  by  a  lo-gram.  If  this  is  too  little,  the  weight 
of  the  crucible  lies  between  10  and  20  grams.  The  5-gram  is 
then  added,  and  if  still  insufficient,  the  2-gram  weight  is  added. 
If  this  is  too  much,  the  2-gram  is  removed  and  a  i-gram  sub- 
stituted. If  this  is  still  too  little,  the  weight  of  the  object  lies 
between  16  and  17  grams.  The  subdivisions  of  the  gram  are  then 
used  in  the  same  systematic  order,  until  the  second  decimal  place 
is  established,  that  is,  until  the  whole  number  of  centigrams  is 
ascertained.  Thus,  suppose  the  weight  is  found  to  be  between 
1674  and  1675,  the  front  sash  is  then  closed,  and  the  system 
brought  into  complete  equilibrium  by  means  of  the  rider,  which  is 
moved  from  division  to  division  along  the  graduated  beam,  until 
the  oscillation  of  the  pointer  is  equal  in  both  directions.  Suppose 
that  when  this  is  the  case,  the  position  of  the  rider  is  midway 

*  It  is  not  necessary,  although  it  is  preferable  that  it  should  be  so,  that  the 
gram  weight  should  be  absolutely  true  to  the  standard  gram,  so  long  as  all  the 
weights  used  bear  the  correct  ratios  to  each  other. 

f  It  is  extremely  rarely  that  the  weights  of  a  reliable  maker,  such,  for 
example,  as  Oertling,  are  ever  found  to  be  untrue. 


Drying  and  weighing  a  Filter.  193 

between  the  third  and  fourth  main  divisions,  the  weight  of  the 
crucible  will  then  be  represented  by  the  number  167435  grams.* 

In  recording  a  weight,  which  should  be  done  immediately  the 
weighing  operation  is  finished,  the  value  of  the  weights  should  first 
be  read  off  from  the  empty  spaces  in  the  box,  and  this  should  then 
be  checked  as  the  weights  are  returned  one  by  one,  beginning  with 
the  highest,  to  their  respective  compartments. 

Except  in  the  case  of  pieces  of  metal,  alloys,  etc.,  substances 
which  have  to  be  weighed  must  never  be  placed  directly  upon  the 
scale-pan,  but  must  be  contained  in  a  suitable  vessel,  such  as  a 
weighing-bottle  or  a  watch-glass,  which  itself  must  be  perfectly 
clean  and  dry. 

The  correct  weight  of  an  object  cannot  be  taken  while  it  is  hot, 
owing  to  the  disturbance  introduced  by  the  warm  air-currents  rising 
from  it.  Substances,  therefore,  which  require  to  be  heated  previously 
to  being  weighed,  must  be  allowed  to  cool  down  to  the  temperature 
of  the  air  in  the  balance-room  (see  Desiccator,  p.  196). 

When  taking  a  weighed  quantity  of  a  substance  for  analysis,  it 
is  usual  to  obtain  its  weight  by  difference.  The 
prepared  powder  is  contained  in  a  light  thin  glass 
stoppered  bottle  (a  convenient  form  of  weighing- 
bottle  is  shown  in  Fig.  17),  and  the  bottle  with  its 
contents  is  weighed. 

A  sufficient  quantity  of  the  powder  is  then  care- 
fully tipped  out  into  the  beaker  or  other  vessel 
selected  in  which  to  operate  upon  the  compound, 
and  the  bottle  is  again  weighed.  The  difference 
between  the  two  weighings  represents  the  weight  of 
the  substance  employed  for  the  analysis. 

Sometimes  the  clean  empty  vessel  in  which  the  substance  is  to 
be  afterwards  treated  is  itself  weighed,  then  a  quantity  of  the  sub- 
stance is  introduced  into  it,  and  the  vessel  again  weighed.  This 
would  be  the  method,  for  instance,  when  the  substance  is  to  be 
operated  upon  in  a  crucible. 

2.  Drying  amdL  weighing  a  Filter. — The  filter-paper  most 
suitable  for  use  in  quantitative  analysis  is  a  specially  prepared 
Rhenish  paper  (Schleicher  and  Schtill,  No.  589),  which  has  been 
treated  with  hydrochloric  and  hydrofluoric  acids  in  order  to  remove 

*  This  method  of  weighing  only  gives  the  absolute  weight  of  any  object 
when  the  bulk  or  volume  of  the  article  is  the  same  as  that  of  the  weights  used 
to  counterpoise  it.  In  ordinary  analytical  operations,  the  difference  between 
the  weight  of  the  air  displaced  by  the  objects  and  that  displaced  by  the  weights 
is  too  insignificant  to  be  considered. 


194         Preliminary  Gravimetric  Manipulations. 

soluble  mineral  matters.  The  paper  is  obtained  already  cut  into 
circles,  the  two  most  useful  sizes  for  general  use  being  respectively 
9  and  n  centimetres  (or,  roughly,  3^  and  4^  inches)  in  diameter. 

In  certain  quantitative  estimations,  as  will  be  seen  later,  it  is 
necessary  to  weigh  a  precipitate  which  has  been  dried  upon  a  filter, 
and  in  such  cases  it  is  needful  to  know  the  weight  of  the  dry  paper 
itself,  in  order  to  deduct  this  from  the  total  weight.  The  process 
of  making  this  simple  determination  will  form  a  useful  exercise  in 
a  number  of  different  operations  which  have  constantly  to  be  made 
in  quantitative  analysis. 

The  apparattis  in  which  to  weigh  the  filter  may  be  either  a  light 
stoppered  tube  or  weighing-bottle  (A,  Fig.  18),  or  a  pair  of  watch- 


glasses  with  ground  edges,  held  together  by  means  of  a  bent  wire 
clip  (B,  Fig.  1 8).  A  duplicate  experiment  may  be  made,  using 
each  of  these  forms  of  apparatus. 

The  bottle  and  the  watch-glasses  are  first  rendered  perfectly 
clean  by  means  of  a  dry  clean  glass-cloth,  and  are  then  made  quite 
dry  by  being  placed  for  about  half  an  hour  in  a  steam-oven ;  the 
bottle  with  its  stopper  out,  and  the  glasses  separated  from  each  other. 

The  steam-oven  consists  of  a  double-walled  copper  vessel  (Fig. 
19),  the  space  between  the  two  walls  being  partly  filled  with  water, 
which  is  maintained  at  the  boiling  temperature  by  means  of  a 
Bunsen  lamp.  In  order  to  keep  the  water  at  a  constant  level,  the 
oven  maybe  provided  with  a  feeding  arrangement,  shown  at /in 
the  figure.  This  consists  of  a  syphon,  the  recurved  end  of  which 
is  constantly  immersed  in  water,  which  is  continuously  flowing  in 
at  the  bottom  of  the  outer  tube,  and  overflowing  by  the  lateral  tube 
into  the  sink.* 

*  In  most  well-appointed  laboratories,  stacks  of  steam-ovens,  heated  by 
steam  from  a  steam-boiler,  form  a  part  of  the  regular  fittings.  The  temperature 
within  a  steam-oven  never  quite  reaches  100°. 


Drying  and  weighing  a  Filter.  195 

The  weighing-bottle  and  watch-glasses  are  placed  upon  the 
floor  of  the  oven  (which  must  be  quite  clean),  or  upon  a  piece  of 
clean  paper  placed  first  upon  the  floor.  At  the  expiration  of  about 
half  an  hour  the  apparatus  is  removed  (the  stopper  being  inserted 
in  the  bottle,  and  the  glasses  placed  within  the  clip),  and  deposited 
in  a  desiccator,  where  it  is  allowed  to  cool. 


FIG.  19. 

The  desiccator  is  an  air-tight  glass  vessel,  in  which  the  enclosed 
air  is  kept  dry  by  means  of  some  convenient  drying  agent,  such  as 
lime,  calcium  chloride,  or  sulphuric  acid.  In  its  most  simple  form, 
it  may  be  an  ordinary  bell-glass  with  a  ground  edge,  standing  upon 
a  thick  piece  of  ground  glass,  over  a  glass  dish  containing  the 
desiccating  agent.  The  ground  edge  of  the  bell  is  greased  with 
resin  cerate,  and  the  objects  to  be  kept  dry  are  supported  upon  a 
perforated  zinc  cover  upon  the  inner  dish  (Fig.  20). 

A  handy  form  of  desiccator,  and  one  more  easily  carried  about, 
is  shown  in  Fig.  21.  The  desiccating  material  is  placed  in  the 
lower  part  of  the  vessel,  and  a  disc  of  perforated  zinc,  resting  upon 
the  shoulders,  serves  as  a  support  for  the  objects  to  be  dried.  The 
whole  is  surmounted  with  a  cover,  ground  to  fit  the  vessel,  which, 


196         Preliminary  Gravimetric  Manipulations. 


as  in  the  other  example,  is  greased  with  resin  cerate.*  When  the 
apparatus  has  cooled  in  the  desiccator,  it  is  next  carefully  weighed, 
the  weight  of  the  empty  tube  and  that  of  the  watch-glasses  being 
separately  noted. 

Two  of  the  filter-papers  are  now  folded  in  the  usual  manner. 
One  of  them  is  then  rolled  loosely,  and  introduced  into  the  weigh- 
ing-bottle ;  the  other  is  carefully  folded,  without  cracking  the  paper 
(because  in  actual  work  it  would  afterwards  be  used  for  filter- 
ing), so  that  it  can  be  placed  between  the  watch-glasses.  Each 
piece  of  apparatus  with  the  enclosed  filter  is  then  weighed.  The 
former  weights  deducted  from  those  now  obtained,  give  the  weights 
of  the  nndried  filters.  The  weighing-bottle  and  the  glasses  (opened 


FIG. 


FIG.  21. 


as  before)  are  then  replaced  in  the  steam-oven,  and  heated  for  half 
an  hour,  after  which  they  are  returned  to  the  desiccator  to  cool, 
and  then  re-weighed.  The  difference  between  the  first  weights 
and  these  last  is  the  weight  of  the  dried  filters. 

By  comparing  the  weights  of  the  undried  and  the  dried  filters, 
the  actual  weight  of  the  moisture  which  such  a  piece  of  filter-paper 
is  capable  of  absorbing  from  the  atmosphere  will  be  ascertained. 

The  tube  and  watch-glasses  should  be  once  more  heated  for 

*  Resin  cerate,  as  supplied  by  the  shops,  is  often  rather  too  stiff.  It  is 
well  in  this  case  to  melt  a  quantity  of  it  in  a  porcelain  dish,  and  add  a  little 
vaseline  to  it.  The  whole  should  then  be  poured  into  a  clean  vessel  to  cool. 
An  ordinary  ointment  pot  with  a  cover  (such  as  the  druggists  use)  is  a  con- 
venient vessel  in  which  to  keep  it.  Dust  and  dirt  should  not  be  allowed  to  get 
into  it,  as  any  small  particles  of  gritty  matter  will  prevent  the  lid  of  the  desic- 
cator from  closing  quite  air-tight. 


Estimation  of  Water  of  Crystallisation.         197 

a  similar  period,  cooled  as  before  in  the  desiccator,  and  again 
weighed.  The  weights  now  obtained  should  be  concordant  with 
the  last ;  if  they  are  not,  the  heating  must  be  repeated  until  two 
consecutive  weighings  are  found  to  agree. 

In  actual  practice  we  do  not  want  to  know  the  weight  of  the 
moisture  expelled  from  the  filter,  or  even  the  weight  of  the  dry  paper 
alone.  It  is  enough  to  know  the  weight  of  the  dry  paper  and  the 
weighing-bottle.  The  process  is  therefore  reduced  to  a  single 
operation.  The  paper  is  placed  in  an  unweighed  bottle,  and  after 
half  an  hour's  drying,  it  is  cooled  in  the  desiccator  and  weighed. 
The  paper  is  then  used  to  receive  the  precipitate  which  has  to  be 
weighed.  Then,  after  proper  treatment,  the  filter  with  the  precipi- 
tate upon  it  is  returned  to  the  same  weighing-bottle  and  heated  in 
the  steam-oven  for  a  suitable  time,  after  which  it  is  allowed  to  cool 
in  the  desiccator  and  weighed ;  the  process  of  heating  and  weighing 
being  repeated  until  two  concordant  weights  are  obtained.  The 
weight  of  the  bottle  plus  paper  deducted  from  the  weight  of  the 
bottle  plus  paper  plus  precipitate,  gives  the  weight  of  the  precipi- 
tate, which  it  is  the  object  of  the  series  of  operations  to  ascertain. 

Most  substances,  especially  when  in  the  state  of  fine  powders, 
are  more  or  less  hygroscopic.  On  exposure  to  the  air,  they  absorb 
and  retain  a  certain  amount  of  moisture.  Some  things — such,  for 
example,  as  copper  oxide,  manganese  dioxide,  charcoal  powder,  etc., 
are  able  to  absorb  in  this  mechanical  way  very  appreciable  quan- 
tities of  moisture.  The  water  so  retained  is  in  almost  all  cases 
entirely  expelled  by  heating  the  substance  for  a  suitable  time  in  the 
steam-oven,  after  which,  in  order  to  preserve  it  in  its  dry  state,  it 
must  be  kept  in  the  desiccator. 

When  the  amount  of  moisture  such  substances  contain  has  to 
be  estimated,  a  convenient  quantity  (from  2  to  10  grams,  depend- 
ing upon  the  amount  of  moisture)  is  weighed  out  into  a  pair  of 
watch-glasses,  and  the  process  conducted  in  the  manner  above 
described  in  the  case  of  the  filter-paper.  If  the  substance  is  one 
which  cannot  withstand  a  temperature  of  100°  without  suffering 
chemical  decomposition,  other  methods  have  to  be  adopted  for 
estimating  the  moisture  it  contains,  which  will  be  described  later. 

3.  Estimation  of  Water  of  Crystallisation.— Most  salts 
containing  water  of  crystallisation  readily  part  with  some  or  all  of 
their  water  when  moderately  heated.  In  those  cases  where  the  salt 
suffers  no  other  change  in  composition  at  the  temperature  necessary 
for  the  expulsion  of  the  water  of  crystallisation,  the  latter  may  be 
estimated  by  ascertaining  the  loss  of  weight  which  results  from 
heating  the  substance.  If  the  salt  loses  its  water  of  crystallisation 


198        Preliminary  Gravimetric  Manipulations. 

at  or  below  the  temperature  of  100°,  the  operation  may  be  con- 
ducted in  the  steam-oven  in  the  manner  described  above  ;  but 
when  higher  temperatures  are  necessary,  an  air-oven  is  employed. 

The  air-oven  (Fig.  22)  is  similar  in  construction  to  the  steam- 
oven,  except  that  it  contains  no  water  between  the  walls.  (The 
oven  is  frequently  only  a  single-walled  copper  chamber,  fitted  with 
a  shelf  or  false  bottom).  It  is  heated  by  a  Bunsen  flame,  and  the 


FIG.  22. 

temperature  of  the  air  within  is  ascertained  by  means  of  a  thermo- 
meter, which  passes  through  the  top  by  means  of  a  cork.  The 
door  of  the  oven  is  usually  furnished  with  a  ventilator,  which  can 
be  opened  or  closed  in  order  to  regulate  the  current  of  air  through 
the  oven  ;  the  escaping  air  passes  out  through  a  second  opening  in 
the  top. 

The  objects  to  be  heated  must  be  placed  upon  the  perforated 
shelf  (or  upon  any  other  convenient  support),  which  is  at  some 


Estimation  of  Water  of  Crystallisation.         199 

considerable  distance  from  the  bottom  of  the  oven,  and  the  bulb 
of  the  thermometer  should  be  as  near  to  the  object  as  possible. 
Articles  to  be  heated  must  not  be  placed  directly  upon  the  floor  of 
the  oven,  as  obviously  the  metal  of  that  part  of  the  oven  will  be 
hotter  than  the  temperature  indicated  by  the  thermometer.  This 
is  more  especially  the  case  when  the  oven  is  a  single-walled 
apparatus. 

Within  reasonable  limits,  it  is  easy  to  maintain  the  temperature 
of  such  an  oven  at  any  desired  point  by  regulating  the  Bunsen 
flame.  As  the  mercury  in  the  thermometer  rises  nearly  to  the  re- 
quired point,  the  lamp  is  gradually  turned  down  until  the  upward 
movement  of  the  mercury  ceases.  If  this  point  is  above  the 
desired  temperature,  and  if  after  a  few  minutes  it  does  not  fall,  the 
flame  should  be  lowered  still  more.  (Thermostats,  or  instruments 
which  serve  as  gas-governors,  can  be  fitted  to  the  air-oven,  so  that 
the  supply  of  gas  necessary  to  maintain  the  oven  at  a  constant 
temperature  is  automatically  regulated.) 

Crystallised  copper  sulphate  may  be  conveniently  used  in  order 
to  illustrate  the  method  of  determining  water  of  crystallisation. 
This  compound  crystallises  with  five  molecules  of  water.  Four  of 
these  it  gives  up  slowly  at  100°,  but  more  rapidly  at  1 10°,  while  the 
remaining  molecule  is  retained  until  the  temperature  has  risen 
above  200°.  In  the  following  exercise,  each  of  these  proportions 
of  water  will  be  separately  determined. 

A  quantity  of  purified  copper  sulphate  is  coarsely  powdered  in  a 
mortar  as  quickly  as  possible,  and  the  powder  placed  in  a  stoppered 
bottle  (a  slight  loss  of  water  will  result  if  the  powdering  process 
involves  a  prolonged  exposure  of  the  salt  to  the  air).  About  two 
grams  of  it  are  then  weighed  out  into  a  pair  of  watch-glasses  and 
clip,  by  first  weighing  the  clean  dry  glasses  and  clip,  and  then 
placing  into  them  a  quantity  of  the  salt  which  is  judged  to  be  about 
two  grams,  and  weighing  again.  The  glasses  are  then  separated, 
and  placed  upon  the  shelf  of  the  air-oven,  the  latter  being  main- 
tained at  a  temperature  between  110°  and  115°. 

At  the  expiration  of  an  hour  the  glasses  are  removed  from  the 
oven,  quickly  put  together  in  the  clip,  and  allowed  to  cool  in  the 
desiccator,  and  then  weighed. 

They  are  then  returned  to  the  oven  and  reheated  for  about 
half  an  hour,  after  which  they  are  again  taken  out  and  cooled 
and  weighed.  If  this  last  weight  differs  appreciably  from  the 
previous  one,  the  process  must  be  repeated  until  the  weight  is 
constant. 


2OO         Preliminary  Gravimetric  Manipulations. 

From  these  data  the  percentage  of  water  lost  by  heating  the 
salt  to  i  io°-i  15°  is  calculated  in  the  manner  shown  in  the  following 
example  : — 

Grams. 

Weight  of  watch-glasses  and  salt       i7'3465 

empty          iS'iSgo 

Weight  of  copper  sulphate  used     2>I57S 

Weight  of  watch-glasses  and  salt  after  first  heating  167545 

.,                ,,                ,,                  second   ,,  167230 

ii                ii                ii                  third      ,,  167230 

Weight  of  glasses  and  original  salt 17-3465 

ii             it             dried        ,,      ...         ...         ...  167230 

Weight  of  water  lost  0-6235 

Since  2*1575  grams  of  salt  lost  0*6235  gram  of  water — 

then ^ =  28-90  =  percentage  of  water  lost  at  iio°-ii5° 

The  calculated  percentage  for  four  molecules  of  water  =  28*91 
therefore  the  error,  as  deficit,  =  o'oi 

The  temperature  of  the  air-oven  is  now  raised  to  between 
210°  and  215°,  and  the  opened  watch-glasses  with  the  salt  are  again 
heated  for  about  an  hour  to  this  higher  temperature.  The  glasses 
are  then  removed,  cooled  in  the  desiccator  and  weighed.  As  in 
the  former  case,  this  is  repeated  until  two  consecutive  weighings 
agree. 

In  the  example  the  following  weights  were  obtained  : — 

Grams 

Weight  of  glasses  and  salt  after  first  heating  to  215°    16-5685 
,,  ,,  ,,  second         ,,  16-5680 

Weight  of  glasses  and  original  salt 17-3465 

,,  ,,  dried        ,,      i6'568o 

Total  weight  of  water  lost 07785 

Since  2*1575  grams  of  salt  have  lost  07785  gram  of  water- 
then          / =  36'o8  =  percentage  of  water  lost  at  215° 

Percentage  calculated  for  five  molecules)  _    , 
of  water  of  crystallisation  \ 

Error,  as  a  deficit,  =  0*06 

In  determining  the  water  of  crystallisation  of  salts  which 
undergo  no  other  chemical  change,  except  the  loss  of  their  water, 
even  when  heated  to  a  high  temperature,  it  is  more  expeditious  to 


Estimation  of  Water  of  Crystallisation.         201 

heat  the  salt  in  a  crucible  by  means  of  a  small  Bimsen  flame. 
Copper  sulphate  could  not  be  treated  in  this  way,  as  at  a  temperature 
somewhat  over  300°  the  anhydrous  salt  suffers  decomposition. 
A  convenient  salt  with  which  to  illustrate  the  method  is  barium 
chloride,  which  crystallises  with  two  molecules  of  water. 

A  crucible  (preferably  of  platinum)  with  its  lid,  is  first  heated 
upon  a  pipeclay  triangle  (as  shown  in  Fig.  23)  for  a  few  minutes, 
and  weighed  after  cooling  in  the  desiccator. 

Into  the  crucible  a  quantity  of  pure,  powered  barium  chloride 
is  introduced,  and  the  weight  again  taken. 

The  weight  of  salt  taken   should  be   between  two  and  three 


FIG.  23. 

grams.*  The  crucible  is  at  first  gently  heated  by  a  small  Bunsen 
flame,  the  lid  being  upon  the  crucible.  The  temperature  is 
gradually  raised  until  the  crucible  attains  a  low  red  heat,  at  which 

*  The  quantity  of  material  taken  for  an  estimation  must  depend  upon 
several  considerations.  Speaking  generally,  the  smaller  the  weight  employed, 
the  more  expeditiously  will  the  analysis  be  made.  At  the  same  time,  errors 
due  to  experiment  will  be  more  magnified,  and  therefore  a  greater  demand  is 
made  upon  the  manipulative  skill  of  the  analyst.  The  relation  between  the 
weight  of  the  substance  which  is  to  be  estimated,  and  the  formula-weight  of 
the  compound  to  be  analysed  has  also  to  be  considered  ;  for  instance,  in  the 
present  case,  owing  to  the  high  atomic  weight  of  barium,  and  the  comparatively 
small  quantity  of  water  present,  the  weight  of  the  latter  constitutes  only  a 
small  part  of  the  total  formula-weight  of  the  crystallised  salt. 


2O2        Preliminary  Gravimetric  Manipulations. 

it  is  maintained  for  about  ten  minutes.     It  is  then  placed  in  the 
desiccator  to  cool,  and  afterwards  weighed. 

The  operations  of  heating  and  weighing  are  repeated  until  the 
weight  is  constant.  The  following  is  an  example  : — 

Weight  of  platinum  crucible  and  salt  47-6275 

»•  »  alone 45*3640 

Weight  of  barium  chloride  taken 2*2635 

Crucible  and  salt  after  first  heating 47-2950 

second 47*2940 

Weight  of  crucible  and  original  salt 47-6275 

,,             »              dried        , 47-2940 

Weight  of  water  lost  o'3335 

Since  2*2635  grams  of  salt  contained  03335  gram  of  water — 

then  '  '  2-26^;  ~'~  =  Percenta£e  of  water  of  crystallisation  =  1473 
Percentage  calculated  on  the  formula  BaCl2,2H2O  =  1475 

Error,  as  deficit,  =  o-o2 

In  cases  where  the  substance  undergoes  other  chemical  changes 
at,  or  near,  the  temperature  at  which  it  parts  with  its  water  of 
crystallisation,  or  with  the  moisture  which  is  mechanically  retained 
by  it,  the  weight  of  the  water  which  it  gives  up  cannot  be 
estimated  by  loss.  The  method  is  therefore  so  modified  that  the 
expelled  water  is  absorbed  by  calcium  chloride,  or  other  suitable 
desiccating  material,  which  is  accurately  weighed  before  and  after 
the  experiment.  The  gain  of  weight  thus  observed  represents  the 
weight  of  water  absorbed.  The  apparatus  for  carrying  out  the 
method  is  shown  in  Fig.  24. 

A  weighed  quantity  of  the  substance  is  heated  in  the  bulb-tube 
B.  By  means  of  an  aspirator,  a  slow  stream  of  air  (dried  by 
passing  through  the  calcium-chloride  tube  A)  is  drawn  through 
the  bulb,  and  the  water-vapour  which  is  being  expelled  is  absorbed 
in  the  weighed  calcium-chloride  tube  D.  The  wash-bottle  W, 
containing  sulphuric  acid,  serves  to  show  the  rate  at  which  air 
is  being  drawn  through  the  apparatus,  and  also  prevents  the 
drying-tube  D  from  absorbing  water-vapour  from  the  aspirating 
arrangement. 

Fitting  up  the  Apparatus. — The  two  drying-tubes  are  first 
perfectly  cleaned  and  dried,  and  provided  with  tight-fitting  corks  * 

"*"  When  ordinary  cones  are  employed  in  fitting  up  apparatus  generally, 
those  having  a  fine,  close  texture  and  free  from  faults  should  alone  be  used. 


Estimation  of  Water  of  Crystallisation.         203 

(either  ordinary  or  caoutchouc)  carrying  tubes  as  shown  in  the 
figure.  They  are  then  nearly  filled  with  dry  granulated  calcium 
chloride,  the  pieces  of  which  should  be  about  the  size  of  grains  of 
wheat.  Powder  must  not  be  put  into  the  tubes,  or  they  will  be 
liable  to  become  blocked ;  therefore,  if  the  stock  of  calcium  chloride 
contains  much  that  is  too  fine  for  the  purpose,  it  may  be  passed 
as  quickly  as  possible  through  a  warm  dry  wire  sieve.  A  little 
plug  of  cotton-wool  is  loosely  pushed  down  upon  the  top  of  the 
calcium  chloride,  and  the  tubes  securely  closed  by  the  corks. 


FIG.  24. 


In  the  case  of  the  tube  D,  a  little  plug  of  cotton- wool  is  first 
introduced,  and  drawn  round  to  the  top  of  the  bulb  by  applying 
suction  to  the  end  of  the  narrow  tube,  in  order  to  prevent  particles  of 
the  calcium  chloride  from  getting  into  the  narrow  tube.  The  cork 
which  closes  the  open  limb  of  this  tube  should  be  cut  off  close  to 
the  glass  with  a  sharp  knife,  and  the  surface  completely  covered 
by  shellac  or  sealing-wax.  This  prevents  the  cork  from  absorbing 
moisture  from  the  air. 


The  cork  should  be  at  first  a  trifle  too  large  for  the  tube  it  is  intended  to  fit. 
It  is  then  carefully  pressed  in  cork-squeezers,  or  rolled  under  the  foot  with  a 
gentle  pressure,  which  will  both  soften  it  and  make  it  a  little  narrower.  The 
hole  should  be  cut  with  a  sharp,  clean  cork-borer  to  the  exact  size  required  to 
tightly  fit  the  tube  which  it  is  intended  to  pass  through  it,  without  having  to 
use  a  round  file  to  enlarge  it. 


204        Preliminary  Gravimetric  Manipulations. 


The  ends  of  the  exit  tubes  are  closed  by  means  of  little  caps, 
made  by  plugging  one  end  of  a  short  piece  of  caoutchouc  tube 
with  a  piece  of  glass  rod.  These  caps  should  remain  on  the  tubes 
except  when  they  are  actually  being  used,  or  being  weighed.  The 
tube  D,  which  has  to  be  weighed,  is  furnished  with  a  loop  of  wire 
(preferably  platinum),  by  means  of  which  it  can  be  suspended  from 

the  arm  of  the  balance. 
The  two  drying-tubes  are 
attached  to  a  hard  glass 
bulb  -  tube  by  means  of 
rubber  or  ordinary  corks 
in  the  manner  shown  in  the 
figure.  If  ordinary  corks 
are  used,  they  should  first 
be  thoroughly  dried  in  the 
steam-oven.  A  simple  form 
of  aspirator  for  small  pres- 
sures is  shown  in  Fig.  25. 
An  ordinary  T-tube  is  at- 
tached to  a  narrow  glass 
tube  about  70  cm.  long, 
which  has  been  bent  into 
a  loop  near  one  end.  When 
a  gentle  stream  of  water 
from  the  tap  is  allowed  to 
enter  through  the  tube  S, 
air  is  drawn  in  through  the 
other  limb  of  the  T-tube 
and  carried  down  the  long 
tube.  By  means  of  the 
clamp  D,  the  rate  at  which 
air  is  drawn  through  the 
apparatus  can  be  regu- 
lated. During  the  experi- 
ment it  should  be  such  that  the  bubbles  pass  through  the  sulphuric 
acid  in  the  little  wash-bottle  W  about  two  per  second. 

To  carry  out  the  Operation. — The  drying-tube  D  is  first  care- 
fully weighed,  without  the  little  caps  which  close  the  ends.  These 
are  instantly  replaced  as  soon  as  the  weight  is  ascertained.  The 
bulb-tube,  after  being  carefully  wiped  and  dried,  is  weighed,  and 
a  suitable  quantity  (2  to  4  grams)  of  the  substance  is  introduced 
into  the  bulb.  This  may  be  accomplished  by  placing  the  powdered 


FIG.  25. 


Estimation  of  Water  of  Crystallisation.         205 

substance  upon  a  strip  of  writing-paper,  folded  into  a  gutter 
sufficiently  narrow  to  pass  into  the  tube  (A,  Fig.  26),  and  depositing 
it  in  the  bulb  by  twisting  the  paper  (B,  Fig.  26).  Before  with- 
drawing the  paper,  it  should  be  turned  over  again  into  its  original 
position,  so  that  on  its  way  out  it  may  not  leave  any  traces  of  the 
powder  in  the  stem  of  the  bulb.  The  bulb  is  now  re-weighed,  the 
increase  being  the  weight  of  the  substance  taken. 

The  two  U-tubes  are  then  attached  to  the  bulb,  the  caps  being 
removed  only  from  the  extremities  which  are  fitted  to  the  bulb- 


FIG.  26. 


tube.  The  aspirator  is  then  set  in  operation,  and  when  the  rate 
at  which  it  draws  air  through  the  acid  in  bottle  W  is  adjusted  to 
the  requirements  of  the  experiment,  the  caps  are  removed  from  the 
drying-tubes,  and  the  bottle  W  is  attached  by  a  narrow  caoutchouc 
tube  to  the  drying-tube  D.  Two  screens,  made  of  asbestos  card- 
board, are  placed  in  the  position  shown  in  the  figure  to  shield  the 
corks  and  the  drying-tubes  from  the  heat  of  the  lamp. 

The  bulb  is  gently  heated  by  means  of  a  small  Bunsen  flame, 
the  temperature  being  gradually  increased  until  the  whole  of  the 
water  is  expelled.  Some  of  the  moisture  will  be  seen  to  condense 

P 


206        Preliminary  Gravimetric  Manipulations. 


upon  the  stem  of  the  bulb-tube,  but  as  the  heated  air  passes  through 
the  apparatus,  this  gradually  vaporises  and  passes  on  into  the 
absorbing-tube  D.  When  the  whole  of  the  water  has  been  expelled, 
tube  D  is  detached  and  re- weighed.  Its  gain  in  weight  represents 
the  amount  of  water  contained  in  the  weighed  amount  of  the 
original  compound. 

When  it  is  desirable  to  heat  the  substance  to  some  definite 
temperature,  the  bulb-tube  is  replaced  by  the  U-tube  (Fig.  27), 
which  can  be  heated  to  the  required  tem- 
perature in  a  bath  of  oil  or  melted  paraffin, 
the  temperature  being  ascertained  by  a 
thermometer  suspended  in  the  bath. 

4.  Determination  of  the  Weight 
of  the  Filter-ash.— In  the  majority  of 
cases,  before  a  precipitate  is  weighed,  it 
is  necessary  that  it  be  subjected  to  a  high 
temperature.  It  has,  in  fact,  often  to  be 
strongly  heated  in  a  crucible,  in  order  to 
convert  it  into  a  product  suitable  for  weigh- 
ing. For  example,  the  precipitate  may  con- 
sist of  a  hydrated  oxide  (perhaps  of  indefinite 
composition)  which  it  is  necessary  to  con- 
vert into  the  oxide  (having  a  known  and 
definite  composition)  by  the  action  of  heat. 
In  such  a  process,  obviously,  the  paper  be- 
comes burnt  up,  and  the  ash  only  is  weighed. 
It  is  therefore  needful  to  know  what  is  the 
weight  of  the  ash  so  produced,  that  it  may 
be  deducted  from  the  total  weight. 
For  this  purpose  a  number  of  the  papers  are  incinerated  one 
after  the  other  in  the  following  manner  : — 

A  clean  crucible,  with  its  lid  (either  a  platinum  or  porcelain 
crucible  may  be  used),  is  supported  upon  a  pipeclay  triangle 
(Fig.  23),  and  heated  by  means  of  a  Bunsen  flame  for  a  few 
minutes.  It  is  then  allowed  to  cool  in  the  desiccator,  and  weighed. 
One  of  the  filters  is  then  folded  as  shown  (a,  Fig.  28),  and  rolled 
into  a  compact  little  bundle  (/;)•  A  piece  of  moderately  stout 
platinum  wire  is  then  wound  a  few  times  round  it,  leaving  a 
sufficient  length  of  wire  to  serve  as  a  handle,*  as  shown  at  c.  The 

*  If,  as  is  sometimes  recommended,  the  wire  is  fused  into  a  glass  tube  for 
a  handle,  the  greatest  care  must  be  taken  to  prevent  any  little  fragments  of 
glass,  which  are  liable  to  chip  off  by  the  bending  of  the  wire,  from  falling  into 
the  crucible. 


FIG.  27. 


Determination  of  a  Filter -ash. 


207 


crucible  is  then  placed  upon  a  clean  sheet  of  white  paper  (or  upon 
a  square  of  glass  standing  on  white  paper),  with  the  lid  by  its  side. 


FIG.  28. 

The  rolled-up  paper,  being  held  over  the  crucible  as  in  Fig.  29,  is 
then  ignited  (set  fire  to)  by  means  of  a  Bunsen  flame.  After  the 
flame  of  the  burning  paper  has  died  out,  the  charred  and  shrunken 
remains  still  smoulder  on  for  some  time,  until  the  carbon  has  burnt 
away  and  a  grey  mass  is  left.  It  may  be  lightly  touched  once  or 
twice  with  the  Bunsen  flame  (it  must  not  be  held  for  any  length  of 
time  in  the  gas-flame),  and  then  allowed  to  drop  into  the  crucible. 


FIG.  29. 

If  the  ash  does  not  easily  fall  away  from  the  wire  as  the  latter  is 
inclined  over  the  open  vessel,  a  gentle  tap  of  the  wire  against  the 
rim  of  the  crucible  will  usually  detach  it. 

This  process  is  repeated  (two  papers  may  be  burnt  at  a  time), 
until  the  ash  from  ten  filters  has  been  obtained.  If  any  minute 
particles  have  fallen  on  to  the  paper  or  glass,  they  are  carefully 
swept  up  into  the  crucible  with  a  small  camePs-hair  brush  or  a 
feather.  The  crucible  is  then  covered,  heated  again  for  a  few 
minutes  as  at  first,  placed  in  the  desiccator  to  cool,  and  weighed. 
The  increase  in  weight  is  the  weight  of  the  ash  from  ten  filters,  and 
by  dividing  the  result  by  ten,  the  average  weight  of  the  ash  of 


2o8        Preliminary  Gravimetric  Manipulations. 

one  is  obtained.  If  the  filters  mentioned  on  p.  193  are  used,  the 
weight  of  the  ash  they  yield  is  indicated  on  the  packets  ;  the  student 
will  therefore  be  able  to  compare  the  result  of  his  determination 
with  the  figure  there  given.  It  will  be  obvious  that  the  larger  the 
number  of  papers  incinerated,  the  more  distributed  will  be  the 
experimental  errors,  and  therefore  the  more  exact  will  be  the  result. 

5.  Preparation  of  Pure  Salts.  —  In  order  to  test  the 
accuracy  of  the  results  of  a  quantitative  analysis,  two  courses  are 
open.  Either  duplicate  analyses  are  made,  so  that  two  concordant 
results  may  be  obtained  ;  or  the  result  of  a  single  analysis  is  com- 
pared with  the  calculated  theoretical  composition  of  the  compound 
under  examination.  It  will  be  obvious  that  it  is  only  when  com- 
pounds of  known  composition  and  of  known  purity  are  being 
analysed  that  the  latter  plan  is  possible  ;  in  all  professional  work 
the  former  method  is  necessarily  adopted.  For  the  student, 
however,  who  is  beginning  the  practical  study  of  quantitative 
methods,  it  is  advantageous  that  his  first  exercises  in  quantitative 
estimations  should  be  made  with  salts  which  are  practically  pure 
compounds.  Purified  salts  (i.e.  salts  which  have  been  so  far 
purified  that  the  amount  of  impurity  present  is  too  small  to 
appreciably  affect  the  analysis)  are  readily  obtained  from  chemical 
manufacturers,  but  it  is  well  for  the  student  himself  to  prepare 
some  of  the  compounds  he  will  use,  in  order  that  he  may  gain 
experience  in  the  general  methods  employed  for  this  purpose. 

i.  Crystallisation. — By  the  process  of  crystallisation  from  a 
suitable  solvent  (repeated  several  times,  if  necessary),  a  salt  may 
be  separated  from  most  soluble  substances  with  which  it  is 
admixed.  The  wider  the  difference  between  the  solubilities  of  the 
substances  to  be  separated,  the  more  easily  is  purification  by  this 
method  effected ;  e.g.  barium  chloride  is  more  readily  freed  from 
calcium  chloride  by  recrystallisation  than  from  potassium  chloride, 
the  solubility  of  the  latter  salt  at  the  ordinary  temperature  being 
only  slightly  different  from  that  of  barium  chloride,  while  calcium 
chloride  is  an  extremely  soluble  salt. 

Certain  salts,  on  the  other  hand,  cannot  be  separated  from  each 
other  by  the  process  of  crystallisation  ;  because  when  the  solution 
containing  them  is  allowed  to  crystallise,  the  two  are  together 
deposited,  not  as  a  mere  mixture  of  crystals  of  the  two  individual 
compounds,  but  in  the  form  of  a  double  salt,  having  a  definite 
composition.  ^ 

The  operation  of  recrystallising  a  salt  is  similar  in  outline  in 
every  case,  with  certain  differences  in  detail. 


Preparation  of  Pure  Salts. 


209 


A  hot  strong  solution  of  the  commercial  salt  is  made  in  water. 
For  this  purpose  a  quantity  of  water,  say  from  200  to  300  c.c.,  is 
heated  in  a  beaker,  either  by  means  of  a  rose  burner,  the  beaker 
being  supported  upon  wire  gauze  as  in  Fig.  30,  or  by  means  of  a 
small  Fletcher's  burner  (Fig.  31),  which  is  a  convenient  lamp  and 
stand  in  one.  A  quantity  of  the  salt  to  be  recrystallised  is  then 
added,  moderate  portions  at  a  time,  until  a  tolerably  strong  solution 
is  obtained.  The  quantity  of  salt  which  will  be  dissolved  will 
depend  obviously  upon  the  solubility  of  the  compound  operated 


FIG.  30. 


FIG.  31. 


upon  ;  but  it  is  sometimes  necessary  to  take  other  considerations 
into  account  in  fixing  a  limit  to  the  strength  of  the  solution  which 
is  made.  For  example,  zinc  sulphate  and  magnesium  sulphate, 
when  deposited  from  solutions  at  a  temperature  below  40°  and  50° 
respectively,  form  crystals  containing  seven  molecules  of  water ; 
but  if  the  crystals  are  deposited  from  solutions  above  these  tempe- 
ratures, they  not  only  have  a  different  form,  but  contain  only  six 
molecules  of  water  of  crystallisation.  Therefore  the  solution  must 
either  be  of  such  a  degree  of  concentration  that  crystals  do  not 
form  until  the  temperature  has  fallen  to  these  degrees,  or  else  the 
temperature  of  the  water  in  which  the  salt  is  dissolved  should  not 
be  higher  than  these  points. 

Unless  the  solution  is  absolutely  clear  and  bright  (which,  with 
the  ordinary  commercial  salts,  will  hardly  ever  be  the  case),  it 


2io        Preliminary  Gravimetric  Manipulations. 

must  be  freed  from  suspended  impurity  by  filtration.  In  order  to 
prevent  the  solution  from  crystallising  in  the  funnel,  and  so  block- 
ing the  filter,  it  is  neces- 
sary to  keep  the  funnel 
warm  during  the  opera- 
tion. This  may  be 
done  by  inserting  the 
funnel  in  a  double- 
walled  metal  jacket, 
made  in  the  shape  of  a 
truncated  cone.  The 
jacket  contains  water 
which  is  heated  by 
means  of  a  Bunsen 
flame  (Fig.  32).  The 
same  result  may  be  ac- 
complished by  winding 

FIG.  32.  a    Piece     of     lead     or 

"  compo  "    pipe    round 
the  funnel,  two  or  three  turns  being  enough,  and  blowing  steam 


F/G.  33. 

through  the  pipe.     The  steam  may  conveniently  be  generated  in 
a  common  tin  can.     The  arrangement  is  shown  in  Fig.  33. 


Preparation  of  Pure  Salts. 


211 


A  still  simpler  plan  for  keeping  the  funnel  warm  is  shown  in 
Fig.  34-  A  short  piece  of  "  compo  "  pipe  is  closed  at  one  end,  and 
bent  at  that  end  into  a  loop  or  ring.  Five  or  six  pinholes  are 
bored  in  the  ring,  and  the  pipe  is  supported  by  means  of  a  clamp, 
so  that  the  ring  is  at  some  little  distance  below  the  body  of  the 
funnel.  The  pipe  is  con- 
nected to  the  gas-supply, 
which  is  regulated  so 
that  little  beads  of  flame 
are  burning  at  the  holes 
in  the  pipe. 

The  beaker  contain- 
ing the  filtered  solution 
is  then  placed  in  a  basin 
or  other  convenient 
vessel  containing  cold 
water,  in  order  to  cool 
the  liquid  as  quickly  as 
possible.  The  solution 
is  also  continually  stirred 
with  a  glass  rod.  In 

this  way  the  salt  is  deposited  in  the  form  of  very  small  crystals, 
whereas  if  allowed  to  cool  slowly  and  to  be  undisturbed,  the  crystals 
would  be  larger  and  better  defined.  In  the  former  condition  the 
salt  is  more  easily  freed  from  the  mother- liquor. 

When  the  solution  is  cold,  the  crystals  are  allowed  to  settle, 
and  the  mother-liquor  decanted  off.  The  residue  is  then  trans- 
ferred to  a  drainer,  which  is  a  circular  disk  of  porcelain  or  glass, 
perforated  with  a  number  of  small  holes,  and  supported  in  a  glass 
funnel.  In  this  way  the  greater  part  of  the  mother-liquor  is 
separated  from  the  crystals.  The  process  is  accelerated  by  partially 
reducing  the  atmospheric  pressure  beneath  the  funnel.  For  this 
purpose  the  stem  of  the  funnel  is  fitted  into  the  mouth  of  a  stout 
conical  flask  by  means  of  a  perforated  caoutchouc  stopper,  and  the 
branch  tube  of  the  flask  is  connected  to  any  suitable  exhausting 
apparatus,  such  as  a  Geissler's  or  Wetzel's  water-pump.  The 
arrangement  is  shown  in  Fig.  35,  where  a  Wetzel's  water-pump  is 
represented,  which  is  one  of  the  best  forms  of  its  kind.  Between 
the  pump  and  the  filtering-flask  an  empty  Woulff's  bottle  should 
intervene,  with  the  tubes  arranged  as  shown.  The  object  of  this 
is  to  intercept  the  water  which  is  sucked  back  from  the  pump 
towards  the  flask,  when  the  pump  is  stopped  while  there  is  still  a 


212        Preliminary  Gravimetric  Manipulations. 

partial  vacuum  in  the  apparatus.  The  water  which  is  thus  driven 
into  the  Woulff's  bottle,  is  automatically  drawn  out  again  when  next 
the  pump  is  set  in  action.  The  rubber  tube  used  for  the  water- 
supply  must  be  canvas-covered  or  canvas-lined,  in  order  to  with- 
stand the  water-pressure  ;  that  employed  for  the  other  connections 
must  be  sufficiently  thick  in  the  walls  to  resist  the  pressure  of  the 
atmosphere,  without  collapsing. 

In  cases  where  the  crystals  are  only  moderately  soluble  in  cold 
water,  they  may  be  rinsed  while  upon  the  drainer  by  pouring  a 
little  cold  water  upon  them.  After  they  have  been  allowed  to 


FIG.  35. 


drain  as  completely  as  possible  in  the  funnel,  they  are  then  spread 
out  upon  a  clean  porous  plate  *  or  tile  to  dry. 

If  the  salt  being  operated  upon  is  one  which  does  not  effloresce 
or  undergo  any  change  on  exposure  to  the  air,  it  may  be  left  until 
quite  dry,  care  being  taken  to  prevent  particles  of  foreign  matter 
from  falling  upon  it.  In  the  case  of  compounds  which  cannot 
safely  be  exposed  so  long,  the  crystals  are  finally  dried  by  gently 
pressing  them  between  folds  of  filter-paper.  When  dry,  the  crystals 
should  be  transferred  to  a  dry  stoppered  bottle. 

After  this  manner,  any  of  the  following  salts  may  for  practice 

*  Porous  plates  are  most  convenient  for  the  purpose,  and  are  cheaper  than 
the  tiles.  They  are  ordinary  plates  which,  on  account  of  some  imperfection 
or  slight  damage,  are  not  worth  carrying  any  further  in  the  process  of  manu- 
facture than  the  "biscuit"  or  unglazed  stage.  A  porous  plate  (or  tile)  must 
not  be  used  a  second  time  (unless  for  another  crystallisation  of  the  same  salt), 
as  it  would  obviously  introduce  impurities  into  the  compounds  being  dried. 


Preparation  of  Double  Salts.  213 

be  recrystallised  :  CuSO4,5H2O  ;  MnSO4,5H2O  ;  ZnSO4,7H2O  ; 
NiSO4,7H2O  ;  MgSO4,7H2O  ;  BaCl2)2H2O  ;  KC1  ;  NH4C1  ; 
K2CrO4  ;  K2Cr2O7. 

Double  salts  are  salts  which  contain  two  metals  and  only 
one  acid  radical.*  They  are  usually  formulated  as  associa- 
tions of  molecules  of  the  two  single  salts  which  enter  into  their 
composition  :  thus,  potassium  aluminium  sulphate  (potassium 
alum),  K2SO4,A12(SO4)3,24.H2O  ;  ammonium  nickel  sulphate, 
(NH4)2SO4,NiSO4,6H2O  ;  potassium  magnesium  sulphate, 
K2SO4,MgSO4,6H2O  ;  and  so  on.f  The  exact  nature  of  the  union 
between  the  molecules  of  the  two  single  salts  is  not  known,  but  it 
is  a  union  which  exhibits  shades  of  stability  just  as  in  the  case  of 
that  which  exists  between  the  atoms  within  the  molecules.  For 
instance,  the  union  between  the  two  sulphates  in  potassium  alum 
is  a  much  closer  or  more  stable  union  than  that  between  the  two 
sulphates  in  potassium  magnesium  sulphate.  When  alum  is  dis- 

*  Although  it  is  true  that  "  double  salts"  contain  two  metals  and  one  acid 
radical,  it  does  not  follow  that  all  salts  which  contain  two  metals  and  one 
acid  radical  belong  to  this  class  of  compounds.  For  instance,  microcosmic  salt, 
HNa(NH4)PO4,  and  ammonium  magnesium  phosphate,  NH4MgPO4,  are 
examples  of  salts  containing  two  metals  (regarding  NH4  as  a  metal)  and  one 
acid  radical,  which  are  not  classed  as  "  double  salts."  Strictly  speaking,  a 
compound  is  only  a  "  double  salt "  when  the  sum  of  the  effective  valencies  of 
the  two  metals  is  not  greater  than  the  basicity  of  the  single  acid  radical.  Thus, 
the  basicity  of  the  acid  radical  (PO4)  is  3.  In  microcosmic  salt  the  sum  of  the 
valencies  of  the  two  metals  Na  and  (NH4),  is  only  2,  while  in  ammonium 
magnesium  phosphate  it  is  3.  On  the  other  hand,  in  such  a  compound  as 
ferrous  ammonium  sulphate,  (NH4)2SO4,  FeSO4,6H2O,  which  is  a  true  double 
salt,  the  sum  of  the  valencies  of  the  metals  present  is  4  (two  monovalent  NH4 
groups  and  one  di-valent  iron),  while  the  single  acid  radical  (SO4)  has  a  basicity 
only  of  2. 

t  Some  chemists  adopt  formulas  which  are  constructed  on  the  model  of  the 
formula  of  a  single  salt ;  thus,  alum  is  represented  as  A12K2(SO4)4>24H2O,  ferrous 
ammonium  sulphate  as  Fe(NH4)2(SO4)2,6H2O ;  and  so  on.  The  use  of  formulae 
of  this  type,  however,  is  more  advantageously  reserved  for  a  special  class  of 
double  salts,  which  differ  from  those  here  mentioned  in  their  behaviour  towards 
reagents  in  a  very  marked  manner.  For  example,  if  barium  chloride  be  added 
to  a  solution  of  either  of  the  double  sulphates  mentioned  above,  all  the  sul- 
phuric acid  radical  is  precipitated  as  barium  sulphate,  while  chlorides  of  both 
metals  are  formed  ;  thus — 

K2SO4,A12(SO4)3  +  4BaCl2  =  4BaSO4  +  sKCl  +  A12C16 

Certain  double  salts,  however — such,  for  instance,  as  the  double  chlorides  of 
platinum  and  the  alkali  metals — do  not  behave  in  this  way.  Thus,  if  silver 
nitrate  be  added  to  a  solution  of  sodium  platinic  chloride,  2NaCl,PtCl4,  the 
whole  of  the  chlorine  is  not  precipitated  as  silver  chloride ;  the  silver  is  only 
capable  of  uniting  with  2  out  of  the  6  atoms  of  chlorine  in  the  compound,  the 
result  being  the  production  of  the  compound  2AgCl,PtCl4  and  two  molecules 
of  NaNO3.  Such  a  compound  as  this  may  with  advantage  be  expressed  by  the 
formula  Na2PtCl6,  and  the  action  of  silver  nitrate  may  be  indicated  by  the 
equation — 

Na2PtCl6  +  2AgNO3  =  2NaNO3  +  Ag2PtCl« 


214        Preliminary  Gravimetric  Manipulations. 

solved  in  water  (without  entering  upon  the  much-debated  question 
as  to  the  condition  of  the  salt  while  in  solution},  the  solution,  on 
evaporation,  again  deposits  the  double  salt  unchanged  ;  but  when 
the  double  potassium  magnesium  sulphate  is  similarly  treated,  the 
salt  undergoes  a  partial  separation,  and  the  crystals  which  are 
deposited  consist  partly  of  the  double  salt  and  partly  of  the  least 
soluble  of  the  two  single  salts,  namely,  the  potassium  sulphate. 

These  facts  have  to  be  kept  in  mind  in  the  preparation  or  puri- 
fication of  double  salts,  for  it  is  obvious  that,  while  the  more  stable 
of  this  class  of  compounds  (such  as  alum)  can  be  purified  by 
recrystallisation,  the  process  is  inadmissible  in  the  case  of  those 
salts  which  readily  separate  into  the  simple  salts  of  which  they  are 
composed.  The  general  method  for  the  preparation  of  double  salts 
is  to  take  the  two  single  salts  in  the  proportion  of  their  formula- 
weights,  make  a  strong  solution  of  each  separately  in  water,  and 
carefully  mix  the  two  liquids.  For  example,  to  prepare  potas- 
sium alum,  the  two  salts,  aluminium  sulphate,  Al2(SO4)3,i8H2O 
(formula- weight  =  612),  and  potassium  sulphate,  K2SO4  (formula 
weight  =  174),  are  taken  in  the  proportions  indicated  by  these 
weights.  (A  suitable  quantity  in  this  case  would  be  one-third 
of  these  numbers  in  grams,  i.e.  204  grams  of  aluminium  sulphate 
and  58  grams  of  potassium  sulphate.)  If  the  salts  are  both  pure 
(so  called)  compounds,  the  weights  employed  may  be  as  near  to 
the  theoretical  proportions  as  will  be  obtained  by  weighing  the 
salts  on  a  rough  balance.  If,  however,  ordinary  commercial  salts 
are  used,  a  slight  excess  of  aluminium  sulphate  should  be  taken,  as 
this  is  likely  to  be  the  less  pure  of  the  two  compounds.  The  two 
salts,  in  separate  beakers,  are  dissolved  in  hot  water,  and  the  solu- 
tions filtered  if  they  contain  any  suspended  impurities.  A  slight 
loss  of  the  salts  may  occur  in  this  operation,  but  by  using 
moderately  small  filters,  and  exercising  care,  the  loss  will  be 
practically  equal  in  both  cases.  The  two  clear  solutions  are  then 
thoroughly  mixed  by  pouring  them  from  one  beaker  to  the  other 
once  or  twice,  carefully  avoiding  any  loss  of  either  liquid.  The 
solution  is  then  cooled,  and  the  crystals  drained  and  dried  in  the 
manner  already  described. 

In  those  cases  where  the  union  between  the  two  component 
salts  in  the  double  compound  is  less  stable,  as  already  mentioned, 
the  more  insoluble  of  the  single  salts  is  liable  to  crystallise  out  in 
greater  or  less  quantity  along  with  the  double  compound.  To 
prevent  this,  a  considerable  excess  of  the  more  soluble  ingredient 
must  be  employed ;  usually  an  excess  equal  to  one-fourth  or  one- 


Preparation  of  Pure  Salts.  215 

third  of  the  formula-weight  is  employed.  Thus,  in  the  case  of  the 
double  potassium  magnesium  sulphate  above  cited,  instead  of  246 
(the  formula-weight  of  MgSO4,7H2O)  and  174  (formula-weight  of 
K2SO4),  the  proportions  used  should  be  246  4-  80  to  174. 

In  special  cases — such,  for  instance,  as  the  double  salts  contain- 
ing ferrous  sulphate  as  one  of  the  components — the  separate  solu- 
tions must  be  made  at  temperatures  lower  than  the  boiling-point, 
in  order  to  guard  against  the  precipitation  of  basic  salts.  Thus,  in 
preparing  ferrous  ammonium  sulphate,  FeSO4,(NH4)2SO4,6H2O, 
the  temperature  of  the  solutions  should  not  be  higher  than  about 
40°  C. 

The  two  salts  are  taken  in  the  proportion  of  their  formula-, 
weights  (FeSO4,7H2O  =  278,  and  (NH4)2SO4  =  132),  and  dis- 
solved in  water  so  as  to  obtain  solutions  as  strong  as  possible  at  a 
temperature  of  about  40°.  The  liquids  are  then  mixed,  and  the 
mixture  acidified  by  the  addition  of  two  or  three  drops  of  dilute 
sulphuric  acid.  The  solution  is  then  cooled  and  crystallised  as  usual. 

For  practice,  such  double  salts  as  the  following  may  be 
prepared  :  NiSO4,(NH4)2SO4,6H2O ;  MnSO4,K2SO4,6H2O  ; 
2(NH4)Cl,CuCl2,2H20. 

Precipitation  by  Change  of  the  Solvent.— Salts  may  be 
thrown  out  of  their  aqueous  solution  by  the  addition  to  the  liquid 
of  some  substance  which  is  miscible  with,  or  soluble  in,  the  water, 
but  which,  when  so  dissolved  or  mixed,  constitutes  a  solution  in 
which  the  salt  previously  dissolved  in  the  water  is  insoluble.  For 
instance,  barium  chloride  is  soluble  in  water  but  is  almost  in- 
soluble in  a  solution  of  hydrochloric  acid  in  water.  If,  therefore, 
hydrochloric  acid  be  added  to  such  an  aqueous  solution,  the  barium 
chloride  is  at  once  thrown  out  of  solution  as  a  fine  crystalline  pre- 
cipitate. 

Again,  ferrous  sulphate  is  soluble  in  water,  but  practically 
insoluble  in  alcohol.  By  adding  alcohol,  therefore,  to  an  aqueous 
solution  of  this  salt,  a  precipitate  of  ferrous  sulphate  in  fine  powder 
is  produced. 

As  illustrations  of  this  method,  the  following  exercises  may  be 
made : — 

(i)  Precipitation  of -Sodium  Chloride  by  Hydrochloric  Acid. — 
About  half  a  litre  of  water  is  saturated  with  common  salt,  and  the 
brine  filtered.*  Into  the  clear  solution,  contained  in  a  beaker,  a 

*  As  salt  is  almost  equally  as  soluble  in  cold  as  in  hot  water,  the  liquid  need 
not  be  heated  ;  and  for  the  same  reason,  should  occasion  arise  where  a  hot 
solution  of  salt  has  to  be  filtered,  there  is  no  necessity  to  use  the  steam-jacket 
for  the  funnel. 


216       Preliminary  Gravimetric  Manipulations. 

stream  of  gaseous  hydrochloric  acid  is  passed.  The  gas  is  gene- 
rated by  acting  upon  common  salt  with  sulphuric  acid  (previously 
mixed  with  water  in  the  proportion  of  1 1  volumes  of  strong  acid 
to  8  of  water,  and  the  mixture  cooled)  in  the  flask  H,  Fig.  36.  In 
order  to  arrest  any  sodium  sulphate  which  might  be  carried  as  fine 
spray  along  with  the  stream  of  gas,  the  latter  is  made  to  bubble 
through  a  strong  solution  of  hydrochloric  acid  contained  in  a  three- 
necked  bottle,  B,  before  being  passed  into  the  brine.  The  tube 
which  delivers  the  gas  into  the  salt  solution  should  be  sufficiently 
wide  not  to  become  stopped  up  with  the  deposited  salt,  and  need 


FIG.  36. 


not  dip  very  far  into  the  solution.  An  ordinary  thistle  funnel,  sup- 
ported as  shown  in  the  figure,  is  convenient. 

The  generating  flask  should  be  provided  with  a  safety  funnel, 
and  the  wash-bottle  B  with  a  tube  which  just  dips  beneath  the 
surface  of  the  liquid.  In  this  way,  should  there  be  any  interruption 
in  the  evolution  of  gas,  there  is  no  possibility  of  the  brine  being 
sucked  back  into  B,  as,  directly  a  reduction  of  pressure  takes  place 
within  the  apparatus,  air  is  drawn  in  through  the  safety  tubes. 

As  the  gas  begins  to  bubble  into  the  brine,  the  salt  almost 
immediately  commences  to  separate  out  as  a  crystalline  powder. 


Preparation  of  Pure  Salts.  217 

When  a  sufficient  quantity  of  the  salt  has  been  precipitated,  the 
crystals  are  allowed  to  settle,  and  the  bulk  of  the  liquid  decanted 
off.  The  salt  is  then  thrown  upon  a  drainer  (which  may  be  con- 
nected to  a  water-pump,  as  on  p.  2 1 2),  and  after  as  much  as  possible 
of  the  mother-liquor  has  drained  away,  the  salt  is  rinsed  two  or 
three  times  with  a  little  strong  hydrochloric  acid  (pure).  It  is  then 
removed  to  a  porous  plate  to  dry,  after  which  it  is  heated  in  a 
porcelain  evaporating-dish,  by  means  of  a  "  rose  "  burner,  until  all 
traces  of  water  and  hydrochloric  acid  are  expelled. 

(2)  Precipitation  of  ferrous  sulphate,  FeSO4,7H2O,  by  alcohol, 
— A  convenient  quantity  of  ferrous  sulphate  (say  100  grams)  is 
dissolved  in  water  at  a  temperature  of  about  40°  C.  The  salt  should 
be  placed  in  a  beaker,  and  a  quantity  of  water  having  a  tempera- 
ture of  about  40°  added  to  it.  The  temperature  of  the  mixture  at 
once  falls,  perhaps  as  much  as  10°,  owing  to  the  absorption  of  heat 
due  to  solution.  The  beaker  is  then  placed  in  a  larger  beaker  or 
other  vessel  containing  water  at  40°,  and  the  solution  constantly 
stirred  until  the  whole  of  the  salt  has  dissolved  and  a  tolerably 
saturated  solution  obtained  at  that  temperature.  The  solution  is 
acidified  with  two  or  three  drops  of  dilute  sulphuric  acid,  and  then 
filtered.  To  prevent  oxidation  by  undue  exposure  to  the  air,  the 
filtering  operation  should  be  accelerated  by  means  of  the  filter- 
pump  *  (p.  2 1 2).  The  perfectly  clear  greenish  solution  thus  obtained 

*  As  ordinary  filter-paper,  when  wet,  is  not  strong  enough,  when  folded  and 
used  in  the  ordinary  way,  to  stand  the  pressure  to  which  it  would  be  exposed  by 
the  use  of  the  filter-pump,  special  devices  must  be  employed  for  its  support. 

One  method  most  suitable  for  such  occasions  as  that  under  consideration, 
where  great  rapidity  is  an  advantage,  is  to  place  in  the  funnel  a  porcelain 
drainer,  and  to  lay  upon  it  a  small  circular  disc  of  ordinary  filter-paper  cut  a 
trifle  larger  (about  2  mm.  all  round)  than  the  drainer.  The  paper  is  moistened 
and  carefully  adjusted  so  as  to  entirely  cover  the  drainer.  [In  some  cases 
where  a  very  high  degree  of  exhaustion  is  necessary,  a  circular  piece  of  muslin 
or  thin  calico  should  be  laid  under  the  paper.] 

Where  the  filter-pump  is  used  for  precipitates  that  are  afterwards  to  be 
dried  and    weighed,   as    in    the 
usual  quantitative  determinations, 
one  of  the  two  following  plans  is 


usually  adopted :  Either  a  spe- 
cially prepared  toughened  paper 
is  used  (Schleicher  and  Schiill, 
No.  575),  or  the  ordinary  paper 
is  supported  at  the  apex  of  the 
cone  by  means  of  a  small  muslin 
or  a  platinum  cone.  When  muslin 
is  used,  a  small  circular  piece  FIG>  37. 

(conveniently  cut  to  the  size  of  a 

penny  piece)  is  folded  as  a  filter  is  folded,  and  placed  in  the  funnel.  The 
folded  filter-paper  is  then  introduced,  being  carefully  pushed  into  its  position 
with  its  apex  within  the  muslin  cone.  The  filter-paper  must  fit  the  funnel 


218        Preliminary  Gravimetric  Manipulations. 

is  poured  into  a  quantity  of  alcohol  (methylated  spirit)  about  equal 
to  the  volume  of  the  ferrous  sulphate  solution.  The  first  effect  will 
be  somewhat  unexpected,  for,  instead  of  any  precipitation  of  crystals, 
a  green  liquid  separates  out  at  the  bottom  of  the  vessel.  On  stir- 
ring the  mixture,  however,  for  a  moment  or  two,  this  liquid  disap- 
pears, and  the  ferrous  sulphate  is  precipitated  as  a  fine  crystalline 
powder.  The  liquor  is  then  decanted  off,  the  crystals  drained  upon 
a  drainer  and  rinsed  with  alcohol.  It  must  then  be  dried  as 
quickly  as  possible  upon  a  porous  plate,  or  by  gentle  pressure 
between  folds  of  blotting-paper. 

Precipitation  by  Double  Decomposition.— The  process 
of  preparing  a  pure  compound  by  this  method  is  practically  identical 
with  that  of  conducting  any  quantitative  precipitation.  It  will  be 
obvious  that,  in  order  to  obtain  a  pure  product,  there  must  be  nothing 
present  which  will  give  a  precipitate  with  the  reagent  to  be  used, 
other  than  the  compound  intended.  For  instance,  suppose  it  is 
desired  to  prepare  pure  barium  carbonate  by  precipitation  from  a 
solution  of  barium  chloride  by  means  of  ammonium  carbonate  ;  the 
barium  chloride  must  itself  be  a  purified  salt,  and  must  be  free 
from  any  metallic  compounds  which  would  give  a  precipitate  with 
ammonium  carbonate.  The  precipitated  compound  is  washed 
with  warm  water  until  perfectly  free  from  any  soluble  salts,  and  is 
finally  dried  in  a  steam-oven  if  it  can  stand  that  temperature  without 
decomposition,  otherwise  it  is  spread  upon  a  porous  plate. 

When  it  is  desired  to  crystallise  a  salt  with  the  object  of  obtain- 
ing large  and  well-defined  crystals,  it  is  necessary  to  modify  the 
methods  so  as  to  cause  the  process  to  take  place  very  slowly,  and, 
as  nearly  as  possible,  with  uniform  regularity.  A  nearly  saturated 
solution  of  the  salt  is  made  at  a  temperature  about  30°  C,  using 
a  considerable  volume  of  water.  The  solution  is  placed  in  a  beaker, 
which  is  then  stood  in  a  larger  beaker  of  warm  water  in  order  to 
render  the  cooling  operation  slower,  and  covered  with  a  clock- 
glass.  When  the  solution  is  cold,  there  will  be  a  small  crop  of 

exactly,  and  not  leave  any  air-spaces  down  the  sides.  If  the  sides  of  the  funnel 
do  not  enclose  an  angle  of  exactly  60°,  the  paper  must  be  folded  so  as  to  make 
it  exactly  fit  the  funnel,  or  the  funnel  must  be  rejected  for  one  of  the  right 
shape.  A  platinum  cone  is  made  by  cutting  from  a  piece  of  foil  a  segment  of 
a  circle  whose  radius  is  15  mm.  (or  \  inch).  The  segment  should  be  nearly 
three-quarters  of  the  entire  circle,  as  shown  in  Fig.  37,  so  that  when  it  is  bent 
together  so  as  to  form  the  cone,  there  will  be  an  overlap  extending  almost  one- 
half  of  the  way  round.  A  few  fine  holes  may  be  pierced  in  the  metal  near  to 
the  apex  of  the  cone,  by  laying  the  flat  piece  of  foil  upon  a  cork,  and  pricking 
it  with  a  fine  needle.  The  cone  is  then  put  into  the  funnel  with  the  paper  just 
as  the  muslin  cone.  Sometimes  a  cone  of  the  toughened  paper  above  men- 
tioned is  used  ;  this  is  less  advantageous,  as  it  practically  introduces  a  double 
filtering  medium,  and  retards  the  process  considerably. 


Preparation  of  Pure  Salts.  219 

crystals  on  the  bottom  of  the  beaker.  The  liquid  is  then  poured 
into  another  beaker,  and  a  few  of  the  best  formed  of  the  crystals 
are  picked  out.  (If  this  first  crystallisation  has  resulted  in  the 
deposition  of  a  solid  cake  or  crust,  and  no  isolated  crystals,  the 
liquid  must  be  returned  to  the  beaker,  and  the  whole  warmed  up 
again,  with  the  addition  of  a  little  more  water,  and  the  solution  set 
to  cool  as  before.)  These  are  dropped  into  the  cold  mother-liquor 
so  as  to  lie  isolated  from  each  other,  and  the  vessel  placed  where 
its  temperature  will  remain  as  nearly  constant  as  possible.  The 
rate  of  evaporation  of  the  liquid  may  be  checked,  if  the  salt  deposits 
too  quickly,  by  partially  covering  the  beaker.  If  circumstances 
make  it  necessary  that  the  operation  be  carried  on  in  a  room  liable 
to  considerable  changes  of  temperature,  it  is  a  good  plan  to  keep 
the  beaker  in  a  larger  vessel  of  water  all  the  time  the  process  is 
going  on.  In  order  that  the  crystals  may  become  equally  developed 
in  every  direction,  so  as  to  approach  as  nearly  as  possible  the 
perfect  form  for  the  particular  compound  being  crystallised,  they 
must  be  turned  at  regular  intervals,  say  every  day,  or  every  second 
day,  according  to  the  rate  of  growth,  so  as  to  rest  upon  a  different 
face.  The  more  slowly  the  crystals  can  be  grown,  other  things 
being  equal,  the  more  perfect  in  form  will  they  be. 

In  cases  where  the  solution  cannot  be  left  exposed  to  the  air  for 
so  long  without  undergoing  some  chemical  change  (for  instance, 
ferrous  sulphate  would  become  oxidised),  the  salt  may  be  slowly 
crystallised,  if  it  is  one  which  is  insoluble  in  alcohol,  by  allowing 
that  liquid  to  very  gradually  diffuse  into  the  aqueous  solution.  To 
take  the  case  of  ferrous  sulphate  as  an  example.  A  strong  cold 
aqueous  solution  of  the  salt  is  placed  in  a  large  wide-mouth 
stoppered  bottle,  which  is  less  than  half  filled  by  it.  A  layer 
of  water  about  2\  c.c.  (i  inch)  deep  is  floated  upon  the  top  of 
the  solution,  and  upon  the  top  of  the  water  a  quantity  of  alcohol 
is  placed,  equal  to  the  volume  of  liquid  already  present.  The 
stopper  is  inserted,  and  the  bottle  placed  where  it  can  remain  for 
some  weeks  undisturbed.  By  slow  degrees  the  alcohol  diffuses 
into  the  aqueous  solution,  and  thereby  causes  the  gradual  deposition 
of  crystals  of  the  salt. 

The  object  of  the  shallow  layer  of  water  dividing  the  alcohol 
from  the  saline  solution  is  to  prevent  the  precipitation  of  the  salt 
at  the  surface  of  contact,  which  would  be  the  case  if  the  alcohol 
floated  immediately  upon  the  solution. 

A  good  plan  for  floating  one  liquid  upon  another  without  causing 
admixture,  is  the  following  : — 

The  heaviest  solution  is  first  poured  into  the  vessel  by  means  of 
a  funnel  with  a  long  stem,  so  as  not  to  splash  the  sides  of  the 
vessel.  A  thin  circular  disc  of  cork,  as  large  as  will  pass  the 
mouth  of  the  bottle,  is  then  lowered  on  to  the  surface  of  the  solution 
by  means  of  a  fine  thread  attached  to  the  centre  of  the  cork  with  a 
little  wire  hook. 

By  means  of  a  stoppered  funnel,  attached  to  a  long  piece  of 


22O        Preliminary  Gravimetric  Manipulations. 

glass  tube  drawn  to  a  narrow  jet  at  the  lower  end,  the  lighter  liquid 
(in  the  above  instance,  the  water}  is  allowed  slowly  to  stream  on  to 
the  cork,  the  funnel  being  gradually  raised  as  the  cork  ascends,  so 
as  to  keep  the  jet  almost  close  to  the  latter,  as  shown  in  Fig.  38. 

When  a  sufficient  layer  of  the  second  liquid  has  been  floated 
upon  the  first,  the  still  less  dense  third  liquid  (the  alcohol  in  the 


FIG.  38. 


example)  is  introduced  in  the  same  way.  As  the  layer  of  the  added 
liquid  gradually  gets  deeper,  and  the  cork  float  is  further  removed 
from  the  line  of  contact,  the  liquid  may  be  run  in  more  quickly, 
without  fear  of  disturbing  the  solution  below.  When  all  has  been 
added  the  cork  is  withdrawn  by  means  of  the  thread,  and  the  bottle 
stoppered.  (The  cork  should  be  put  into  hot  water  to  soak,  if  it  is 
intended  to  use  it  again  for  a  similar  purpose.) 


SECTION   II. 
TYPICAL  GRAVIMETRIC  ESTIMATIONS  OF  METALS. 

IN  order  to  make  a  gravimetric  estimation  of  a  metal,  the  latter  is 
either  obtained  from  the  compound  in  the  elemental  form,  and  the 
metal  itself  weighed  ;  or  it  is  caused  to  enter  into  some  convenient 
combination,  yielding  a  compound  of  known  and  definite  com- 
position, and  from  the  weight  of  the  compound  so  produced  the 
quantity  of  metal  is  ascertained  by  calculation.  This  latter  mode 
is  of  more  general  application  than  the  first,  as  many  of  the  metals 
are  not  easily  reduced  from  their  compounds.  Of  late  years,  how- 
ever, electrolytic  methods  of  analysis  have  been  greatly  developed 
and  perfected,  so  that  a  large  number  of  metals,  which  formerly 
could  only  be  determined  in  combination,  are  now  deposited  in  the 
form  of  metal,  and  so  weighed.  Typical  examples  of  this  process 
of  analysis  will  be  given  in  a  separate  section. 

The  forms  of  combination  which  are  most  frequently  had  recourse 
to,  and  which  are  applicable  to  various  metals,  are  practically  three 
in  number,  namely,  the  oxide,  sulphide,  and  sulphate. 

When  weighed  in  the  form  of  an  oxide,  the  metal  is  first  pre- 
cipitated either  as  a  hydroxide  (as,  for  example,  in  the  case  of 
Al,  Fe,  Cr,  Ni,  Cu,  etc.)  or  as  a  carbonate  (as  in  the  case  of  Bi,  Mn, 
Zn,  etc.).  In  either  case  the  precipitated  compound  is  converted 
by  heat  into  the  oxide,  which  is  then  weighed. 

When  estimated  as  a  sulphide,  the  metal  is  precipitated  in  this 
form  of  combination  either  by  sulphuretted  hydrogen  or  ammonium 
sulphide.  Examples  of  metals  which  may  be  estimated  in  this  way 
are  seen  in  mercury,  cadmium,  antimony,  and  arsenic.  Metals 
which  are  to  be  weighed  as  sulphates,  are  either  precipitated 
from  solution  in  this  form,  as  in  the  case  of  lead,  barium,  and 
strontium,  or  the  sulphate  is  obtained  by  evaporating  a  suitable 
compound  of  the  metal  (e.g.  the  chloride)  with  sulphuric  acid, 
whereby  it  becomes  converted  into  the  sulphate.  This  method  is 
sometimes  employed  for  the  estimation  of  the  alkali  metals. 

Q 


222  Typical  Gravimetric  Estimations. 

Besides  these  methods,  each  of  which  is  applicable  for  the 
estimation  of  several  metals,  there  are  special  forms  of  combination 
which  are  available  for  special  cases.  Thus,  for  example,  silver  is 
precipitated  and  weighed  as  chloride  (or  bromide)  ;  magnesium 
is  precipitated  as  ammonium  magnesium  phosphate,  which  is  then 
converted  into  magnesium  pyrophosphate  by  the  action  of  heat, 
and  weighed  in  this  form  ;  and  so  on. 

Descriptions  will  be  given  in  this  section  of  a  selected  number  of 
typical  examples,  illustrating  the  most  important  methods  of  gravi- 
metric estimations,  and  forming  a  suitable  course  for  practice  in 
the  various  analytical  operations. 

Estimation  of  Aluminium. 

Epitome  of  Process. — The  metal  is  precipitated  from  the 
solution  of  its  compounds  as  aluminium  hydroxide  by  means  of 
ammonia.*  The  hydroxide  is  afterwards  converted  by  heat  into 
the  oxide,  A12O3,  in  which  form  the  metal  is  weighed. 

Employ  potassium  alum,  K2SO4,A12(S04)3,24H2O,  taking  about 
2  grams.f 

A  stoppered  weighing-bottle  containing  powdered  alum,  which 
has  been  purified  by  re-crystallisation,  is  first  weighed  ;  after  which 
about  2  grams  of  the  salt  are  carefully  transferred  without  loss  to 
a  beaker,  \  conveniently  of  about  550  c.c.  (nearly  20  ozs.)  capacity. 
(The  stopper  should  be  removed  from  the  bottle  while  the  lattei 
is  held  immediately  over  the  beaker.  A  sufficient  quantity  of  the 
salt  is  then  tipped  out  and  the  stopper  again  replaced,  care  being 
taken  that  no  particles  are  left  adhering  to  the  lip  of  the  bottle, 
by  giving  the  latter  a  few  taps  with  the  finger,)  The  beaker  is 

*  The  precipitation  of  metallic  hydroxides  by  alkalies  is  not  complete  in  the 
presence  of  certain  organic  substances,  such  as  tartaric  or  citric  acids,  sugar, 
etc.  In  cases  where  such  non-volatile  organic  matter  is  present,  these  must 
first  be  decomposed  by  adding  to  the  solution  sodium  carbonate  and  potassium 
nitrate,  and  evaporating  to  -dryness  in  a  platinum  dish,  upon  a  steam-bath. 
The  residue  is  then  fused,  and  afterwards  extracted  with  dilute  hydrochloric 
acid.  The  solution  is  then  filtered,  and  the  precipitation  conducted  as  above 
described. 

t  See  note  on  p.  201.  In  this  case,  the  formula  weight  of  alum  being  948, 
while  that  of  ALOs  is  only  102,  the  final  product  that  is  to  be  weighed  amounts 
to  less  than  one-ninth  of  the  weight  of  the  original  salt.  Hence  a  fair  quantity 
must  be  operated  upon.  Two  grams  of  the  salt  will  only  furnish  about  o'ar 
grams  of  A12O3.  When  estimating  the  (SO4)  in  alum  the  case  is  different,  for 
the  weight  of  the  BaSO4  yielded  is  nearly  equal  to  the  original  weight  of  the 
alum  itself. 

t  In  special  cases,  where  exceptional  accuracy  is  required,  a  vessel  of 
platinum,  nickel,  or  porcelain  is  substituted  for  the  beaker,  as  alkalies  exert  a 
slight  solvent  action  upon  the  glass.  With  Jena  glass  this  action  is  extremely 
slight. 


Estimation  of  Aluminium.  223 

immediately  covered  with  a  clock-glass.    The  bottle  is  then  weighed 
again,  and  the  difference  gives  the  weight  of  alum  taken. 

The  salt  is  then  dissolved  in  such  a  quantity  of  water  that  the 
beaker  is  about  one-fourth  part  filled  with  liquid,*  and  the  process 
of  solution  is  aided  by  gently  warming  the  mixture. 

For  this  purpose  the  beaker  may  be  supported  and  heated  by 
either  of  the  methods  shown  on  p.  209,  Figs.  30  and  31  ;  or  it 
may  be  placed  upon  an  iron  plate  heated  by  a  Bunsen,  as  in  Fig.  39. 


FIG.  39. 


This  is  a  specially  convenient  arrangement  when  several  operations 
are  being  conducted  simultaneously.  The  heat  can  be  moderated 
by  placing  the  vessel  nearer  to,  or  farther  from  the  flame.  A  piece 
of  asbestos  cloth  laid  upon  the  top  of  the  plate  prevents  any  risk  of 
the  fracture  of  glass  vessels  by  placing  them  suddenly  upon  the 
rough  iron.  It  is  important  that  the  various  pieces  of  apparatus 
used  should  bear  a  suitable  relation  to  each  other  as  regards  their 
size.  For  example,  the  clock-glass  employed  as  a  cover  should  not 
be  much  larger  than  the  mouth  of  the  beaker ;  and  when  a  flame 
is  used  directly  under  the  vessel,  care  must  be  taken  that  it  does 
not  extend  beyond  the  bottom  of  the  beaker.  Such  an  arrangement 

*  In  all  the  following  examples,  the  size  of  the  beaker,  and  the  quantity  of 
water  employed  for  the  solution  of  the  salt,  may  be  the  same  as  those  here 
described,  unless  special  directions  to  the  contrary  are  given.  It  ought  not  to 
be  necessary  to  repeat  at  this  stage,  what  was  stated  on  pp.  3  and  116,  that 
in  all  analytical  operations  distilled  water  must  be  exclusively  used. 


224 


Typical  Gravimetric  Estimations. 


FIG.  40. 


of  apparatus  as  shown  in  Fig.  40,  and  which  is  not  unfrequently 
seen  in  the  laboratory,  is  almost  sure  to  end  in  disaster.  The  heat 
from  the  lamp  either  fractures  the  clock-glass,  fragments  of  which 
fall  into  the  solution,  or  the  beaker  itself 
becomes  unduly  heated  upon  its  sides,  so 
that  the  first  movement  of  the  liquid 
causes  it  to  crack.  The  beaker  should 
be  kept  covered  by  the  clock-glass  as 
much  as  possible  throughout  the  entire 
operation. 

A  quantity  of  ammonium  chloride 
solution  is  added,*  equal  to  about  a  fourth 
of  the  volume  of  the  liquid  already  present, 
and  then  ammonia  in  the  least  possible 
excess  beyond  that  which  is  required  for 
complete  precipitation. 

The  mixture  is  stirred  with  a  glass  rod 
with  rounded  ends,  the  rod  being  of  such 
a  length  that  it  will  just  stand  in  the 
beaker  while  the  cover  is  upon  it.  A  short 
piece  of  clean  narrow  caoutchouc  tube 
should  be  previously  fitted  upon  the  end 
of  the  rod.  This  not  only  prevents  the  beaker  from  becoming 
scratched  by  the  rod,  but  it  is  useful  in  the  subsequent  operation  of 
removing  particles  of  the  precipitate  which  may  adhere  to  the  sides 
of  the  beaker. 

The  liquid  is  then  boiled  until  the  greater  part  of  the  ammonia 
is  expelled,  and  the  issuing  steam  possesses  only  a  slight  smell  of 
the  reagent. 

The  precipitate  is  then  allowed  to  settle  ;  the  clock-glass  is  re- 
moved, and  the  under  side  rinsed. into  the  beaker  by  means  of  a 
fine  jet  of  hot  water  from  a  wash-bottle,  in  case  any  particles  may 
have  been  thrown  up  during  the  operation  of  boiling.  The  clear 
liquid  is  then  decanted  off  through  a  filter,!  the  weight  of  whose  ash 
is  .known  (see  p.  206). 

When  as  much  of  the  liquid  as  possible  has  been  poured  off 

*  In  the  presence  of  ammonium  chloride,  aluminium  hydroxide  is  insoluble 
in  ammonia. 

f  The  operation  of  filtering  must  be  carried  out  with  all  the  precautions 
against  loss  described  in  the  introduction  to  qualitative  analysis,  p.  2.  It 
sometimes  happens,  owing  to  the  particular  shape  of  the  lip  of  a  beaker,  that, 
even  when  the  liquid  is  carefully  poured  down  against  a  glass  rod,  a  little  of  it 
creeps  back  under  the  edge.  This  may  be  prevented  by  slightly  greasing  the 
under  side  of  the  lip  by  means  of  a  little  touch  of  vaseline  or  resin  cerate  upon 
the  finger. 


Estimation  of  Aluminium. 


225 


without  disturbing  the  precipitate,  the  beaker  is  half  rilled  up  with 
boiling  water,  the  mixture  well  stirred,  and,  after  being  allowed  to 
settle,  is  decanted  as  before  through  the  same  filter.  This  process 
is  repeated  a  second  time,  after  which  the  precipitate  itself  is  trans- 
ferred to  the  filter.  Every  trace  is  removed  from  the  sides  of  the 
beaker  by  the  aid  of  the  glass  rod,  and  the  beaker  is  finally  rinsed 
into  the  funnel  by  means  of  a  jet  of  hot  water  from  the  wash-bottle 
in  the  manner  shown  in  Fig.  41,  the  movable  nozzle  of  the  bottle 
being  directed  round  the  beaker  by  the  fore-finger. 

The  final  washing  of  the  precipitate  is  now  made  upon  the  filter 
by  means  of  the  jet  of  hot  water  from  the  wash-bottle,  washing 
downwards  from  the 
upper  edge  of  the 
paper,  but  never  quite 
filling  the  paper  cone 
with  water,  and  allow- 
ing each  washing  to 
drain  completely 
through  before  adding 
fresh  water.*  The 
operation  is  continued 
until  the  precipitate  is 
entirely  freed  from  all 
soluble  salts,  which  in 
this  case  consist  of 
sulphates  of  potassium 
and  ammonium.  To 
ascertain  when  this  is 
the  case,  a  few  drops 
of  the  filtrate  are  col- 
lected in  a  test-tube, 
and  a  little  barium  chloride  added.  If  no  precipitate  forms  in  a  few 
minutes  after  warming  the  mixture,  the  washing  may  be  considered 
complete.  The  last  washing  should  aim  at  collecting  the  precipitate 
together  into  the  cone  of  the  filter. 

The  filter  and  precipitate  are  next  dried  in  a  steam-oven.  For 
this  purpose,  the  funnel  containing  the  filter  is  covered  with  a  piece 

*  When  a  filter-pump  is  being  used,  however,  this  rule  does  not  apply.  In 
this  case  it  is  important  not  to  let  the  filter  run  empty,  but  to  constantly  add 
liquid  before  the  former  quantity  has  completely  run  through,  otherwise 
channels  or  fissures  are  liable  to  be  formed  in  the  solid  material  upon  the  filter 
through  which  subsequent  wash-waters  run  without  properly  washing  the 
precipitate. 


FIG.  41. 


226 


Typical  Gravimetric  Estimations. 


of  filter-paper,  and  either  stood  up  in  the  corner  of  the  oven,  or 
placed  in  a  conical  tin  support.  In  order  to  securely  cover  the 
funnel,  a  common  circular  filter-paper,  somewhat  larger  in  diameter 
than  the  top  of  the  funnel,  is  held  upon  the  latter  with  the  left  hand, 

while  the  projecting 
paper  is  folded  down 
all  round  with  the  other 
hand  in  the  manner 
shown  in  Fig.  42. 

While  the  precipi- 
tate is  drying,  a  plati- 
num crucible  is  heated 
to  redness  for  a  few 
moments,  cooled  in 
the  desiccator,  and 
then  weighed.  The 
dried  filter  containing 
the  aluminium  hydrox- 
ide is  then  removed 
from  the  funnel,  care- 
fully folded  up,  and  placed  in  the  weighed  crucible.  The  crucible, 
supported  upon  a  clean  pipeclay  triangle,  is  first  very  gently 
heated,  the  lid  being  placed  slightly  on  one  side  in  order  to  allow 
the  gaseous  products  of  the  combustion  of  the  paper  to  escape,  and 
also  to  admit  the  necessary  supply  of  air  for  the  combustion.  When 
the  active  combustion  of  the  paper  has  subsided,  the  temperature 
is  gradually  raised,  and  when  it  has  reached  a  red  heat,  the  lid  of 
the  crucible  may  be  removed.*  The  vessel  is  maintained  at  a  bright 
*  It  is  very  necessary  that  the  tips  of  the  crucible-tongs  should  be  perfectly 


FIG.  42. 


>ssn  ^yir 


FIG.  43. 

clean,  and  after  being  in  contact  with  the  hot  crucible  or  lid,  the  tongs  should 
never  be  placed  on  the.  table  with  the  tips  downward.    The  first  impulse  of  almost 


Estimation  of  Aluminium. 


227 


red  heat  by  means  of  a  good   Bunsen   flame   or  a  blowpipe  for 
about  ten  minutes.     The  lid  is  then  replaced,  and  the  crucible 

every  one  is  to  hold  and  use  crucible-tongs  in  the  manner  shown  in  Fig.  43, 
with  the  almost  inevitable  result  that  when  putting  them  down,  the  hot  points 


come  into  contact  with  the 
wooden  table  (Fig.  44),  and 
are  liable  to  pick  up  matter 
which  afterwards  may  be  trans- 
ferred to  the  crucible.  Tongs 
should  always  be  held  in  ex- 
actly the  reverse  way,  with 
the  tips  pointing  in  the  direc- 
tion of  the  back  of  the  hand, 
as  in  Fig.  45.  They  are  then 
naturally  placed  upon  the  table 
with  the  tips  pointing  upwards 
(Fig.  46).  The  habit  of  mani- 
pulating tongs  in  this  way 
should  be  cultivated  from  the 
very  beginning. 


FIG.  46. 


228  Typical  Gravimetric  Estimations. 

transferred  to  the  desiccator  to  cool,  after  which  it  is  weighed. 
It  is  reheated  for  a  similar  time  and  again  weighed,  the  operation 
being  repeated  until  the  weight  is  practically  constant. 

[When  a  filter-pump  is  employed  for  accelerating  the  filtration 
of  the  aluminium  hydroxide,  the  precipitate  may  be  partially  dried 
by  allowing  the  pump  to  continue  in  operation  for  fifteen  minutes 
or  so  after  all  the  liquid  has  passed  through.  The  filter  may  then 
be  folded  up,  and  at  once  transferred  to  the  crucible,  in  which  it  is 
completely  dried  by  the  cautious  application  of  heat.  To  avoid 
spitting,  the  crucible  should  be  supported  rather  obliquely  in  the 
triangle,  and  a  small  flame  applied  to  the  extreme  upper  part,  so 
that  the  heat  may  be  conducted  very  gently  to  the  wet  materials.] 

From  the  data  obtained  by  the  analysis,  the  percentage  of 
aluminium  is  calculated  in  the  following  manner  :  — 

Estimation  of  Aluminium  in  Potash  Alum,  K2SO4,- 
A12(S04)3,24H20.- 

Weigbing-bottle  and  salt  (first  weight)    ...         ......     16-9485  grams. 

,,  ,,  ,,       (second  weight)  ......     14-8870      t, 

Weight  of  salt  taken  ......      2-0615      ,, 

Weight  of  crucible  +  filter-ash  +  A12O3  .........     25-6857 

Weight  of  crucible  alone    .........  25-4635 

Weight  of  ash*       ............  0-0002 

---    25-4637 

WTeight  of  A12O3  ............      0-2220 

(Formula-weight  of  A12O3  =  102 
i,    _A12      -     54 

therefore  weight  of  aluminium  ^      54  x  0*2220  _ 
in  o*2220gram  Al2O3t  ;  102 

o*ii7in  x  100 
hence  the  percentage  of  aluminium  found  =—  -  -T  --  —  5'7° 


*  If  the  filter  employed  is  that  mentioned  on  p.  193,  having  an  ash  equal  to 
0-00017,  it  may  be  ignored  altogether,  as  it  does  not  affect  the  second  place  of 
decimals  in  the  final  result. 

f  The  ratio  102  :  54  is  equal  to  i  :  0-5294,  therefore  the  weight  of  Al  in  any 
given  weight  of  A12O3  is  at  once  obtained  by  multiplying  the  latter  by  the 
factor  0-5294.  Thus,  in  the  above  example,  0-2220X0-5294  —  0*117526.  At 
the  commencement  of  each  estimation  in  this  section,  the  factor  will  be  found, 
which,  when  multiplied  by  the  weight  of  the  compound  that  is  actually  weighed, 
will  give  the  weight  of  metal  it  contains.  The  use  of  factors,  especially  by  the 
student,  is  always  attended  with  the  danger  of  obscuring  the  rationale  of  the 
calculation,  and  reducing  it  to  a  mechanical  or  "rule  of  thumb"  operation. 
To  prevent  this  as  far  as  possible,  the  ratio  of  the  formula-weights  of  the  com- 
pound weighed,  and  the  metal  to  be  estimated  are  given,  so  as  to  keep  present 
to  the  mind  the  true  significance  of  the  factor. 


Estimation  of  Chromium.  229 

^Formula-weight  of  K2SO4,A12(SO4)S,24H2O  =  948 
\  Formula-weight  of  A12  =    54 

hence  the  theoretical  percentage  of  |  _  54  x  100  _ 

~~ 


_ 
aluminium  in  alum  /  ~~       948 

Chromium. 

Epitome  of  Process.  —  The  chromium  is  precipitated  from 
solutions  of  chromic  compounds  in  the  form  of  chromic  hydroxide 
by  means  of  ammonia.  The  hydroxide  is  converted  by  heat  into 
chromic  oxide,  Cr2O3,  in  which  form  the  element  is  weighed. 

Chromates  are  first  reduced  to  the  "  chromic  "  state  by  means 
of  sulphur  dioxide.  Factor  — 

(Cr2O3)  152  :  (Cr2)  104  =  i  :  0-68421 

Estimation  of  Chromium  in  Chrome  Alum,  K2SO4,- 
Cr2(SO4)3,24H2O.—  Employ  about  i'5  gram.  The  process  is  con- 
ducted exactly  as  in  the  case  of  aluminium,  except  that  the  addition 
of  ammonium  chloride  may  be  omitted.  The  filter  may  either  be 
burnt  in  the  crucible  along  with  the  precipitate,  or  it  may  be  in- 
cinerated separately,  as  described  below  in  the  case  of  iron. 
Chromium  oxide  is  not  reduced  during  the  process. 

Formula  weight  of  chrome  alum  =  998.  Theoretical  percentage 
of  Cr  =  10-42. 

Estimation  of  Chromium  in  Potassium  Bichromate, 
K2Cr2O7.  —  Take  about  075  gram.  The  purified  salt  is  weighed  out 
into  a  beaker  (see  notes,  pp.  222,  223),  and  dissolved  in  water  without 
being  warmed.  A  gentle  stream  of  sulphur  dioxide  *  is  passed 
through  the  liquid  until  the  chromium  is  wholly  reduced  to  the 
"  chromic  "  state.  The  reduction  is  complete  when  the  solution, 
on  being  gently  moved  against  the  sides  of  the  beaker,  shows  no 
trace  of  the  yellow  colour,  but  has  a  clear  green  tint.  The  delivery 
tube  is  rinsed  into  the  beaker  with  water  from  the  wash-bottle,  and 
removed.  The  solution  is  then  warmed  to  expel  the  excess  of 
sulphurous  acid,  and  is  afterwards  precipitated  with  ammonia  as 
in  the  above  example. 

Formula-weight  of  potassium  dichromate  =  294.  Theoretical 
percentage  of  Cr  =  35  '37. 

*  The  sulphur  dioxide  is  most  conveniently  derived  from  a  "syphon"  of 
the  liquefied  gas.  In  the  absence  of  such  a  supply,  the  gas  must  be  generated 
from  sulphuric  acid  and  copper.  In  this  case  the  sulphur  dioxide  must  be 
washed  by  being  passed  through  water  before  being  delivered  into  the  solution 
of  the  chromate. 


230  Typical  Gravimetric  Estimations. 

Iron. 

Epitome  of  Process. — The  iron  is  precipitated  from  solutions 
of  ferric  compounds  as  ferric  hydroxide  by  means  of  ammonia. 
The  hydroxide  is  converted  by  heat  into  ferric  oxide,  Fe^O3,  in 
which  form  the  metal  is  weighed.  The  filter  is  burnt  apart  from 
the  precipitate.  The  iron  in  ferrous  compounds  is  first  oxidised 
into  the  "  ferric  "  state.  Factor — 

(Fe2O3)  160  :  (Fe2)  112  =  i  :  0700 

Estimation  of  Iron  in  Ferrous  Ammonium  Sulphate, 

(NH4)2SO4,FeSO4,6H2O.— Employ  about  1-5  gram.  The  purified 
salt  is  weighed  out  into  a  beaker  (notes,  pp.  222,  223),  and  dissolved 
in  water  with  the  addition  of  one  or  two  drops  of  dilute  sulphuric 
acid.  The  solution  is  warmed,  and  a  little  strong  nitric  acid 
added,  sufficient  to  oxidise  the  whole  of  the  iron.  The  mixture  is 
heated  nearly  to  the  boiling-point  for  a  short  time  in  a  covered 
beaker.  Each  drop  of  nitric  acid  as  it  is  added  produces  a  brown 
coloration,  owing  to  the  absorption  of  the  disengaged  nitric  oxide 
by  the  ferrous  salt ;  but  when  the  oxidation  is  complete,  the  addi- 
tion of  a  drop  of  the  acid  produces  no  visible  effect.  By  adding, 
therefore,  a  few  drops  of  nitric  acid  after  its  addition  has  ceased  to 
cause  any  coloration,  sufficient  will  have  been  introduced  to  com- 
plete the  oxidation.  As  a  confirmation,  the  smallest  drop  of  the 
solution  is  placed  upon  a  piece  of  white  porcelain  (such  as  a 
crucible  lid)  by  means  of  a  glass  rod,  and  the  drop  is  touched  with 
another  glass  rod  which  has  been  dipped  in  a  solution  of  potassium 
ferricyanide.  If  the  ferrous  iron  has  been  completely  oxidised,  no 
blue  coloration  will  result.  [The  ferricyanide  solution  must  be 
freshly  made,  by  dissolving  a  small  crystal  of  the  salt  in  water.] 

The  ferric  hydroxide  is  then  precipitated  by  the  addition  of  a 
slight  excess  of  ammonia,  and  the  mixture  boiled  until  the  steam 
scarcely  smells  of  ammonia.  The  precipitate  is  washed  with  hot 
water  by  decantation  two  or  three  times,  and  finally  washed  upon 
the  filter  until  the  wash-water  is  entirely  free  from  sulphates  (com- 
pare Aluminium,  p.  225). 

The  precipitate  is  then  thoroughly  dried  in  the  steam-oven. 

When  perfectly  dry,  the  filter  is  withdrawn  from  the  funnel,  and 
as  much  of  the  precipitate  as  possible  is  detached  from  the  paper 
by  gently  squeezing  the  cone  together,  and  transferred  to  a  platinum 
crucible,  which  is  placed  upon  a  sheet  of  glazed  paper  (see  p.  207). 
By  flattening  the  paper  cone  and  gently  rubbing  one  side  against 
the  other,  the  remaining  adhering  particles  may  be  detached.  The 
paper  is  then  folded  in  the  manner  described  on  p.  207,  bound  up 


Estimation  of  Calcium.  231 

in  a  platinum  wire,  and  incinerated  as  completely  as  possible.  The 
ash  is  then  shaken  off  the  wire  into  the  crucible,  and  any  particles 
which  may  have  fallen  upon  the  glazed  paper  are  also  carefully 
swept  into  the  crucible  by  means  of  a  feather.  The  crucible  is 
then  heated  for  about  ten  minutes  to  a  bright  red  heat  by  means  of 
a  Bunsen  flame,  after  which  it  is  placed  to  cool  in  the  desiccator 
and  weighed  ;  the  heating  and  weighing  being  repeated  until  no 
further  loss  of  weight  results. 

The  minute  quantity  of  ferric  oxide  which  is  left  upon  the  filter, 
and  which  becomes  reduced  to  metallic  iron  during  incineration,  is 
again  oxidised  during  the  process  of  heating  ;  and,  although  it 
may  not  pass  into  the  same  oxide,  Fe2O3,  but  probably  into  Fe3O4, 
the  total  quantity  is  so  small,  and  the  ratio  Fe3O4  :  (Fe3)  is  so 
nearly  equal  to  Fe2O3  :  (Fe2),  that  the  difference  is  quite  outside  the 
experimental  error. 

Formula  weight  of  (NH4)2SO4,FeSO4,6H2,0  =  392.  Theoreti- 
cal percentage  of  Fe  =  14*28. 

Calcium. 

Epitome  of  Process. — The  calcium  is  precipitated  in  the  form 
of  calcium  oxalate,  CaC2O4,  by  means  of  ammonium  oxalate  in  the 
presence  of  ammonia.  The  calcium  oxalate  is  then  converted  by 
a  gentle  heat  into  calcium  carbonate,  CaCO3,  in  which  form  the 
metal  is  weighed  ;  or  it  is  strongly  heated  until  it  is  entirely 
changed  into  calcium  oxide,  and  weighed  in  this  form. 

As  an  exercise  in  manipulation,  the  oxalate  may  be  converted 
first  into  the  carbonate  and  weighed,  and  afterwards,  by  the  appli- 
cation of  a  stronger  heat,  be  changed  into  the  oxide  and  weighed 
again.  Factors — 

(CaCO3)  loo  :  (Ca)  40  =  I  :  0*4000 
(CaO)  56  :  (Ca)  40  =  i  :  071428 

Estimation  of  Calcium  in  Calcium  Carbonate,  CaCO3.— 
Employ  about  0-5  to  075  gram.  Pure  precipitated  calcium  carbonate 
is  weighed  out  into  a  beaker,  and  dissolved  in  a  very  little  dilute 
hydrochloric  acid.  The  clock-glass  cover  should  be  pushed  a  little 
to  one  side,  and  the  diluted  acid  poured  gradually  down  the  side  of 
the  beaker.  When  the  carbonate  is  completely  dissolved,  the  cover 
should  be  rinsed  into  the  beaker  by  means  of  the  wash-bottle,  and 
water  added  until  about  the  usual  volume  is  present.  Ammonia  is 
then  added  until  the  solution  smells  distinctly  of  the  reagent,  and 
the  liquid  is  heated  to  boiling.  The  calcium  oxalate  is  then  pre- 
cipitated by  the  addition  of  a  slight  excess  of  a  warm  strong  solution 
of  ammonium  oxalate,  to  which  a  little  ammonia  has  been  added. 


232 


Typical  Gravimetric  Estimations. 


The  mixture  is  boiled  for  a  few  minutes,  and  allowed  to  settle.* 
The  clear  liquid  is  then  decanted  off  through  a  filter,  without  dis- 
turbing the  precipitate.  It  is  washed  three  or  four  times  by  decan- 
tation  with  boiling  water,  allowing  it  to  settle  thoroughly  each  time. 
The  small  quantity  of  the  precipitate  which  is  inevitably  conveyed 
to  the  filter  with  the  washings,  will  so  far  fill  up  the  pores  of  the 
paper  that,  when,  after  the  third  or  fourth  wash,  the  precipitate  itself 
is  transferred  to  the  filter,  the  filtrate  will  be  perfectly  clear.  The 
precipitate  is  washed  with  warm  water  until  the  wash-water  is  free 
from  ammonium  chloride,  as  indicated  by  the  absence  of  any 
milkiness  on  the  addition  of  silver  nitrate  after  acidifying  with 
nitric  acid. 

When  washing  this  precipitate,  a  fine  jet  of  water  should  be 
employed,  and  it  should  be  made  to  impinge  first  upon  the  glass 
funnel  above  the  edge  of  the  paper,  and  then  gradually  directed 
down  towards  the  precipitate. 

After  being  dried  in  the  steam-oven,  the  precipitate  is  trans- 
ferred as  completely  as  possible  to  a  platinum  crucible.  The  filter 

is  incinerated  in  a  plati- 
num wire  coil,  and  the  ash 
deposited  in  the  crucible. 
The  calcium  oxalate  is 
then  converted  by  a  gentle 
heat  into  the  carbonate, 
the  temperature  being 
carefully  regulated  so  that 
it  never  reaches  a  visible 
redness  even  at  the  bottom 
of  the  crucible.  A  con- 
venient plan,  in  order  to 
avoid  unduly  heating  any 
spot  of  the  crucible,  is  to 
place  it  upon  a  triangle 
arranged  above  a  piece  of 
wire  gauze  upon  a  tripod,  as  shown  in  Fig.  47.  The  gauze  is 
heated  to  redness  by  means  of  a  rose  burner,  the  heat  being  thus 
radiated  to  the  crucible,  which  must  be  heated  in  this  manner  for 

*  Calcium  oxalate  is  a  precipitate  requiring  some  special  care  in  its  mani- 
pulation. Under  ordinary  circumstances  it  is  somewhat  slow  to  settle,  showing 
an  inclination  to  creep  up  the  wet  sides  of  the  beaker,  and  it  also  passes  through 
the  filter  very  easily.  By  conducting  the  precipitation  in  the  manner  here 
described,  the  compound  is  obtained  in  a  more  coherent  form,  and  with  a  little 
care  will  be  perfectly  retained  by  the  filter. 


FIG.  47- 


Estimation  of  Calcium.  233 

about  fifteen  minutes ;  it  is  then  cooled  in  the  desiccator  and 
weighed. 

In  case  any  of  the  calcium  carbonate  has  become  converted 
into  the  oxide  in  spite  of  care  in  the  heating,  the  residue  is 
moistened  with  a  few  drops  of  a  strong  solution  of  ammonium 
carbonate,  which  is  then  evaporated  to  dryness  by  placing  the 
crucible  in  the  steam-oven.  When  dry,  it  is  reheated  for  a  few 
minutes  over  the  heated  gauze,  and  again  weighed. 

If  this  process  results  in  any  increase  in  the  weight,  it  shows 
that  in  the  first  heating  some  of  the  carbonate  had  been  decom- 
posed ;  the  series  of  operations  should,  therefore,  be  repeated  until 
the  weight  remains  constant. 

In  this  particular  instance  the  form  of  combination  in  which 
the  metal  to  be  estimated  is  weighed,  happens  to  be  the  same  as 
that  in  which  it  exists  in  the  original  compound,  namely,  the 
carbonate;  hence  the  weight  of  the  product  obtained  will  be  the 
same  as  that  of  the  substance  taken  for  analysis. 

The  calcium  carbonate  is  now  converted  entirely  into  calcium 
oxide  by  heating  the  crucible  to  a  strong  red  heat  by  means  of  a 
Bunsen  flame  for  about  ten  minutes,  and  finishing  off  with  a  blow- 
pipe flame.  The  crucible  is  placed  in  the  desiccator  to  cool,  and 
is  then  weighed.  It  is  afterwards  reheated  with  the  blowpipe  for 
a  few  minutes  and  weighed  again,  the  process  being  repeated  until 
no  further  loss  of  weight  results. 

Magnesium. 

Epitome  of  Process. — The  magnesium  is  precipitated  by 
sodium  phosphate,  in  the  presence  of  ammonium  chloride  and 
ammonia,  as  ammonium  magnesium  phosphate,  NH4MgPO4,6H2O. 
This  is  afterwards  converted  by  heat  into  magnesium  pyrophosphate, 
Mg2P2O7,  in  which  form  the  metal  is  weighed.  Factor — 

(Mg2P2O7)  222  :  (Mg2)  48  =  i  :  0-2162 
Estimation  of  Magnesium  in  Magnesium  Sulphate, 

MgSO4,7H2O. — Take  about  75  gram.  The  recrystallised  salt  is 
weighed  out  into  a  beaker,  and  dissolved  in  a  rather  smaller  bulk 
of  water  than  usual  (about  60  to  80  c.c.).  A  quantity  of  ammonium 
chloride  solution  is  added,  equal  to  about  one-half  the  volume  of 
the  solution  of  magnesium  sulphate,  and  then  ammonia  in  mode- 
rate excess.*  To  this  is  added  a  solution  of  hydrogen  disodium 

*  Should  the  addition  of  ammonia  cause  precipitation,  more  ammonium 
chloride  must  be  introduced  until  this  precipitate  is  redissolved.  Excess  of 
ammonia,  over  and  above  that  which  is  required  for  the  formation  of  the 
ammonium  magnesium  phosphate,  is  necessary  to  ensure  complete  precipitation, 
as  the  phosphate  is  slightly  soluble  in  water,  but  insoluble  in  the  presence  of 
free  ammonia. 


234  Typical  Gravimetric  Estimations. 

phosphate,  in  quantity  in  excess  of  that  required  for  complete  precipi- 
tation of  the  ammonium  magnesium  phosphate.  The  mixture  must 
'be  well  stirred  with  a  rubber-tipped  glass  rod.  The  precipitate  is 
a  crystalline  compound,  which  requires  several  hours  for  its  com- 
plete separation.  Its  formation  is  accelerated  by  stirring  or  shaking, 
but  care  is  necessary  to  avoid  pressing  or  rubbing  the  glass  of  the 
stirring-rod  against  the  sides  of  the  beaker,  as  this  would  cause  tho 
deposition  of  minute  crystals  upon  the  surface  of  the  vessel,  which 
are  very  difficult  to  detach.  The  covered  beaker  should  be  put 
aside  for  a  few  hours.* 

The  liquid  is  filtered,  and  the  precipitate  washed  Avith  dilute 
ammonia  (water  3  parts  by  volume,  strong  ammonia  sp.  gr.  '880 
i  part)  until  the  filtrate  is  perfectly  free  from  chloride,  as  indicated 
by  the  absence  of  any  milkiness  when  tested  with  silver  nitrate 
after  acidification  with  nitric  acid. 

The  precipitate  is  then  dried  in  the  steam-oven. 

The  dry  precipitate  is  transferred  to  a  platinum  crucible,  and 
the  filter  incinerated  in  a  platinum  coil.  Care  should  be  taken  not 
to  overheat  the  filter  during  this  operation,  whereby  the  adhering  pre- 
cipitate becomes  fused,  and  thereby  renders  the  complete  combustion 
of  the  paper  extremely  difficult.!  After  the  ash  has  been  placed  in 
the  crucible,  the  latter  is  gently  heated  by  means  of  a  small  Bunsen 
flame.  The  crucible  must  be  covered  during  this  operation,  and 
the  heat  must  be  applied  cautiously,  as,  during  the  conversion  of 
the  compound  into  the  pyrophosphate,  ammonia  and  water  are 
eliminated,  and,  in  escaping  too  rapidly,  these  would  cause  loss  of 
the  precipitate.  When  the  smell  of  ammonia  can  no  longer  be 
perceived,  the  temperature  is  gradually  raised  to  a  bright  red  heat, 
finishing  off  with  a  blowpipe  flame. 

The  crucible  is  then  removed  to  the  desiccator  to  cool,  and 
afterwards  weighed,  the  heating  process  being  repeated  until  the 
weight  is  constant. 

Formula  weight  of  MgSO4,7H2O  =  246.  Theoretical  percentage 
of  Mg  =  9756. 

*  If  the  entire  operation  be  carried  out  in  a  500  c.c.  wide-mouth  stoppered 
bottle,  and  the  mixture  be  briskly  shaken  for  a  few  minutes,  the  precipitation 
will  be  complete  in  a  very  short  time.  The  disadvantage  of  this  plan  is  the 
greater  difficulty  of  transferring  the  precipitate  to  the  filter.  When  the  stopper 
is  removed,  it  must  be  carefully  rinsed  with  water. 

f  The  incineration  may  be  made  by  the  alternative  method  described  for 
magnesium  arsenate,  p.  254. 


Estimation  of  Copper,  235 

Copper. 

I.  As  COPPER  OXIDE. 

Epitome  of  Process.* — The  copper  is  precipitated  from  solu- 
tions of  its  salts  in  the  form  of  copper  hydroxide  by  means  of  potas- 
sium hydroxide.  The  copper  hydroxide  is  subsequently  converted 
by  heat  into  copper  oxide,  in  which  form  the  metal  is  weighed. 
The  filter  is  burnt  apart  from  the  oxide,  and  treated  in  a  special 
manner.  Factor — 

(CuO)  79-3  :  (Cu)  63-3  =  i  :  079823 

Estimation  of  Copper  in  Copper  Sulphate,  CuSO4,5H2O. 
— Employ  about  I  gram.  The  recrystallised  salt  is  weighed  out 
into  a  beaker,f  and  dissolved  in  water  with  the  aid  of  heat.  The 
solution  is  raised  to  the  boiling-point,  and  potassium  hydroxide 
added  in  small  quantities  with  constant  stirring,  until  precipitation 
is  complete.  When  this  is  the  case,  the  liquid  will  appear  colour- 
less as  the  precipitate  settles,  and  a  drop  of  the  clear  liquid  placed 
on  turmeric  paper  by  means  of  a  glass  rod,  will  show  that  the 
solution  is  alkaline. 

The  mixture  is  gently  boiled  for  a  few  minutes,  and  then  allowed 
to  settle. 

The  precipitate  is  washed  two  or  three  times  by  decantation 
with  boiling  water,  and  then  transferred  to  the  filter,  when  the 
washing  is  continued  with  hot  water  from  the  wash-bottle  until  the 
filtrate  is  perfectly  free  from  potassium  sulphate,  as  indicated  by 
barium  chloride.  The  funnel  is  then  covered,  and  placed  to  dry  in 
the  steam-oven. 

The  dried  precipitate  is  transferred  to  a  platinum  crucible,  and 
as  much  as  possible  of  that  which  adheres  to  the  paper  is  detached 
by  the  process  of  gently  rubbing  the  sides  of  the  filter  together. 

The  paper  is  folded  up  for  incineration,  so  that  that  portion 
which  is  soiled  with  the  precipitate  shall  be  in  the  interior  of  the 
roll,  and  shall  not  come  in  contact  with  the  platinum  wire. 

By  giving  the  crucible  a  gentle  tap,  the  precipitate  is  made  to 
take  up  a  position  rather  to  one  side,  and  so  leave  a  clear  space, 
upon  which  the  ash  may  be  deposited. 

After  the  ash  has  been  dropped  into  the  crucible,  a  single  drop 

*  See  note  *,  p.  222,  with  reference  to  the  presence  of  organic  matter.  In 
the  case  of  copper  hydroxide,  precipitation  is  also  incomplete  if  much  alkali 
nitrate  is  present.  In  such  cases  the  copper  may  be  estimated  as  sulphide 
(p.  236). 

f  See  note  J,  p.  222,  which  applies  still  more  in  the  case  of  caustic  potash 
or  soda. 


236  Typical  Gravimetric  Estimations. 

of  strong  nitric  acid  is  allowed  to  fall  upon  it  from  a  pipette.  Any 
particles  of  reduced  copper  are  thus  converted  into  the  nitrate, 
which,  in  the  subsequent  heating  operation  is  decomposed,  leaving 
copper  oxide. 

The  crucible  is  first  gently  warmed  to  expel  the  excess  of  nitric 
acid,  after  which  it  is  raised  to  a  red  heat,  and  maintained  at  this 
temperature  for  about  ten  minutes.  It  is  then  allowed  to  cool  in 
the  desiccator,  and  weighed.  The  heating  is  repeated  until  the 
weight  is  practically  constant. 

Formula  weight  of  CuSO4,sH2O  =  249-3.  Theoretical  per- 
centage of  Cu  =  25-35. 

II.  As  COPPER  SULPHIDE. 

Epitome  of  Process.  —  The  copper  is  precipitated  from 
neutral  or  acid  solutions  as  cupric  sulphide,  CuS,  by  means  of 
sulphuretted  hydrogen.  The  cupric  sulphide  is  then  converted 
by  heat  into  cuprous  sulphide,  Cu2S,  in  which  form  the  metal  is 
weighed.  Factor — 

(Cu2S)  158-6  :  (Cu2)  126-6  =  i  :  079823 

Estimation  of  Copper  in  Copper  Sulphate,  CuSO4,5H2O. 
—Take  about  I  gram.  The  weighed-out  salt  is  dissolved  in  hot 
water,  in  quantity  sufficient  to  about  half  fill  the  beaker,  and  a  few 
cubic  centimetres  of  strong  hydrochloric  acid  are  added.  A  stream 
of  sulphuretted  hydrogen  is  then  passed  through  the  solution,  until 
precipitation  is  complete. 

The  precipitate  is  at  once  filtered,  the  funnel  being  covered  with 
a  clock-glass  to  prevent  atmospheric  oxidation  of  the  copper 
sulphide  into  sulphate,  which  would  pass  into  solution  and  be  lost.* 
The  filter  should  be  continuously  replenished,  and  not  allowed  to 
run  dry.  The  precipitate  is  washed  down  into  the  apex  of  the 
filter  with  warm  sulphuretted  hydrogen  water,  and  as  in  this 
instance  the  solution  contains  no  soluble  salts,  two  or  three  rinses 
with  the  warm  sulphuretted  hydrogen  water  will  be  sufficient  to 
wash  out  the  acid  present.f  The  filter  is  then  put  to  dry  in  the 
steam-oven. 

The  dry  precipitate  is  detached  from  the  filter  and  transferred  to 
a  Rose's  crucible  \  (previously  heated  and  weighed).  The  paper  is 

*  The  filtrate  may  be  tested  to  see  if  any  copper  is  thus  escaping  into  the 
solution  by  passing  sulphuretted  hydrogen  through  a  portion  of  it. 

f  By  double  decomposition  sulphuric  acid  will  have  been  formed,  and  if 
this  is  not  washed  out  the  paper  will  become  charred  when  it  is  dried  in  the 
steam-oven. 

%  This  is  a  porcelain  crucible  provided  with  a  perforated  lid,  through  which 
a  porcelain  tube  can  be  passed,  by  means  of  which  the  contents  of  the  crucible 


Estimation  of  Copper. 


237 


then  incinerated  in  the  usual  way,  and  added  to  the  precipitate. 
A  little  powdered  sulphur  (purified  by  redistillation)  is  introduced 
into  the  crucible,  which  is  then  covered  with  its  perforated  lid. 
A  gentle  stream  of  dry 
hydrogen  is  allowed  to  flow 
through  the  bent  porcelain 
tube  (Fig.  48),  which  is 
then  inserted  into  the  hole 
in  the  lid  until  the  flange 
upon  the  tube  rests  upon 
the  lid.  The  crucible  is  first 
gently  heated  until  the  sul- 
phur is  nearly  all  burnt  off, 
when  the  temperature  is 
raised  to  a  bright-red  heat 
and  maintained  at  this  point 
for  about  ten  minutes.  The 
apparatus  is  then  allowed 
to  cool,  with  the  hydrogen 
still  passing  through  it,  until 
it  is  nearly  cold,  when  it  is 
removed  to  the  desiccator, 
and  after  becoming  quite 
cold  is  weighed. 

The  operation  is  repeated,  with  the  addition  of  a  little  more 
sulphur,  until  the  weight  is  practically  constant. 

Cuprous  thiocyanate,  Cu2(CyS)2.  Copper  may  also  be  pre- 
cipitated as  cuprous  thiocyanate,  which  can  be  afterwards  converted 
into  cuprous  sulphide,  Cu2S,  by  the  action  of  heat  in  an  atmosphere 
of  hydrogen  in  a  Rose's  crucible  ;  or  it  may  be  weighed  on  the 
filter  as  cuprous  thiocyanate. 

The  solution  of  the  copper  salt,  which  should  be  slightly  acidified 
with  hydrochloric  acid,  is  first  saturated  with  sulphur  dioxide  *  (see 

can  be  heated  in  an  atmosphere  of  hydrogen.  In  the  absence  of  such  an 
apparatus,  an  ordinary  porcelain  crucible  may  be  used  ;  and  a  common  clay 
tobacco-pipe  (of  such  a  size  that  the  mouth  of  the  inverted  bowl  will  just  pass 
into  the  crucible)  may  be  substituted  for  the  perforated  flange  and  tube. 
Hydrogen  from  a  cylinder  of  compressed  gas  is  most  convenient.  If  the  gas 
has  to  be  generated  from  sulphuric  acid  and  zinc,  it  must  be  dried  before 
entering  the  crucible,  by  being  passed  through  strong  sulphuric  acid.  By 
arranging  the  flow  of  gas  through  the  tube  before  introducing  it  into  the 
crucible,  there  is  no  risk  of  projecting  the  contents  out  by  too  suddenly 
admitting  the  gas.  Although  not  so  advantageous,  coal-gas  may  be  employed 
instead  of  hydrogen. 

*  Or  a  solution  containing  equal  weights  of  hydrogen  ammonium  sulphite 
and  ammonium  thiocyanate  may  be  used  to  precipitate  the  copper  salt. 

R 


FIG.  48. 


238  Typical  Gravimetric  Estimations. 

note  on  p.  229).  A  solution  of  ammonium  thiocyanate  is  then  added 
until  precipitation  is  complete,  and  the  mixture  heated  nearly  to 
the  boiling-point  for  a  few  minutes.  The  precipitate,  which  is 
coloured  when  it  is  first  formed,  quickly  becomes  white,  and  settles 
very  readily.  It  is  washed  with  warm  water  by  decantation  two  or 
three  times,  and  then  transferred  to  the  filter,  where  it  is  washed 
until  the  wash -water  is  free  from  the  acid  of  the  original  copper 
salt  (eg,  sulphuric  acid  if  copper  sulphate  is  employed  for  the 
estimation).  The  precipitate  is  then  dried  in  the  steam-oven.  The 
dry  precipitate  is  then  treated  as  described  above  in  the  case  of 
copper  sulphide,  and  finally  weighed  as  cuprous  sulphide. 

As  an  alternative  method,  the  precipitated  cuprous  thiocyanate 
may  be  filtered  upon  a  weighed  filter-paper,  and  dried  in  a  hot-air 
oven  at  a  temperature  of  110°  to  115°,  and  finally  weighed  as 
cuprous  thiocyanate,  the  heating  being  repeated  until  the  weight 
is  constant.  The  compound  retains  water  very  persistently. 
Factor— 

(Cu2(CyS)2)  242-6  :  (Cu2)  126-6  =  i  :  o  52184 

Silver. 

Epitome  of  Process.— Silver  is  precipitated  from  solutions 
of  its  salts  in  the  form  of  silver  chloride,  AgCl,  by  means  of 
hydrochloric  acid,  and  weighed  in  this  form.*  The  filter  is  inciner- 
ated apart  from  the  precipitate,  with  special  precautions.  Factor — 

(AgCl)  H3'16  '  (Ag)  107-66  =  i  :  075202 

Estimation  of  Silver  in  Silver  Nitrate,  AgNO3.— Use 
about  0-5  gram.  The  purified  salt  is  weighed  out  into  a  beaker,  and 
dissolved  in  cold  water,  and  the  solution  acidified  with  a  little 
nitric  acid.  Dilute  hydrochloric  acid  is  then  gradually  added,  with 
continued  stirring,  until  precipitation  is  complete.  This  point  is 
easily  determined,  as  the  precipitated  silver  chloride  coagulates  and 
settles  very  quickly,  so  that  it  is  easy  to  see  when  the  addition  of 
acid  produces  no  further  precipitation.  The  mixture  is  heated 
nearly  to  the  boiling-point,  and  then  allowed  to  settle.  The  pre- 
cipitate is  washed  two  or  three  times  by  decantation  with  boiling 
water  acidified  with  nitric  acid,  after  which  it  is  transferred  to  the 
filter,  and  washed  with  hot  water  until  the  filtrate  is  free  from 
hydrochloric  acid,  as  indicated  by  the  addition  of  silver  nitrate. 
It  is  then  placed  in  the  steam-oven  to  dry. 

*  Silver  is  sometimes  precipitated  and  weighed  as  silver  bromide.  The 
precipitation  is  effected  by  means  of  a  solution  of  ammonium  bromide,  as 
hydrobromic  acid  is  not  a  usual  laboratory  reagent.  For  the  same  reason  the 
filter  ash  is  treated  as  described  in  the  alternative  method,  No.  2,  for  the 
chloride,  given  above.  Owing  to  the  greater  weight  of  bromide,  the  experi- 
mental error  is  more  equally  divided  between  the  silver  and  the  halogen  than 
in  the  case  of  the  chloride. 


Estimation  of  Silver.  239 

The  dry  precipitate  is  detached  from  the  paper  as  thoroughly  as 
possible  and  transferred  to  a  porcelain  crucible,  which  with  its  lid 
has  been  previously  heated  and  weighed.  The  paper  is  then  folded 
up  with  the  soiled  parts  innermost,  and  incinerated  in  a  platinum 
coil.  The  greatest  care  must  be  taken  to  arrange  the  folding  of  the 
paper  and  the  binding  of  the  wire  round  it  in  such  a  way  that 
neither  the  silver  chloride  nor  the  reduced  silver  comes  in  contact 
with  the  heated  wire  ;  otherwise  the  platinum  will  become  alloyed 
with  silver,  and  the  analysis  will  be  vitiated. 

The  filter  ash  is  deposited  in  the  inverted  lid  of  the  crucible, 
where  it  is  moistened  with  a  single  drop  of  strong  nitric  acid, 
allowed  to  fall  upon  it  by  means  of  a  pipette.  The  reduced  silver  is 
thereby  converted  into  the  nitrate.  With  a  similar  pipette  a  single 
drop  of  hydrochloric  acid  is  added,  which  reproduces  silver  chloride. 
The  crucible  lid  is  then  cautiously  heated  upon  a  pipeclay  triangle 
by  means  of  a  small  flame  placed  at  a  considerable  distance  below 
it,  until  the  acids  are  completely  evaporated. 

The  crucible  is  then  gently  heated  until  the  precipitate  just 
begins  to  melt,  when  it  is  removed  along  with  the  lid  to  the 
desiccator,  and  weighed  when  cold.* 

Formula  weight  of  silver  nitrate  =  i69'66  theoretical  percentage 
of  Ag  =  63-45. 

Alternative  Methods  of  Treating  thePilter.f—  i.  After 
the  precipitate  has  been  transferred  to  the  crucible,  the  latter  is 
heated  until  the  silver  chloride  just  begins  to  fuse.  The  filter  is 
then  rolled  up  and  incinerated  as  described  above,  and* the  ash 
allowed  to  fall  into  the  crucible.  It  is  then  moistened  with  the 
acids  as  already  explained,  and  the  crucible  again  heated ;  at  first 
very  gently,  to  expel  the  excess  of  acid,  afterwards  to  the  point  at 
which  the  silver  chloride  begins  to  melt.  The  crucible  is  then 
cooled  and  weighed. 

2.  The  precipitate  is  transferred  to  the  crucible,  which  is  then 
heated  until  the  chloride  begins  to  fuse.  The  crucible  is  cooled  in 
the  desiccator  and  weighed. 

The  filter  is  then  burnt  as  before,  and  the  ash  deposited  in  the 
crucible,  which  is  once  more  gently  heated,  cooled,  and  weighed. 
The  increase  in  weight  is  made  up  of  the  weight  of  the  ash  plus  the 

*  The  fused  silver  chloride  may  afterwards  be  removed  from  the  crucible  by 
adding  a  little  dilute  sulphuric  acid,  and  introducing  a  fragment  or  two  of 
granulated  zinc.  The  chloride  is  thereby  reduced  to  metallic  silver,  and  the 
mass  becomes  detached  from  the  porcelain. 

f  A  third  alternative  process  is  similar  to  that  given  for  lead,  Method  A, 
p.  241. 


240  Typical  Gravimetric  Estimations. 

weight  of  a  small  quantity  of  metallic  silver.  This  weight  of  silver 
is  added  to  the  amount  which,  by  calculation,  is  found  to  be 
present  in  the  silver  chloride  already  weighed.  The  following 
example  will  make  clear  the  process  of  calculating  the  result : — 

Weight  of  silver  nitrate  taken        ...         ...         ...     0*292  gram 

Crucible -f  AgCl        21 7265  grams 

Crucible  alone  21-4850      ,, 

Weight  of  AgCl        0-2415 

Multiplying  this  result  by  the  factor,  the  weight  of  metallic  silver 
it  contains  is  obtained — 

0*2415  x  075202  =  0*181612  gram  of  Ag 

Crucible  -f  AgCl  +  ash  +  reduced  silver 21 73067 

Crucible  +  AgCl        217265 

Ash       0-00017 

21  -72667 

Weight  of  reduced  silver          0-004 

Therefore  total  silver=  0*181612  +  0*004  =  0*185612  gram 

.  0*185612  x  loo 
and  -  -2— =  63*5  per  cent,  of  silver 

Lead. 

I.  As  LEAD  OXIDE. 

Epitome  of  Process. — The  lead  is  precipitated  from  solutions 
of  its  salts  in  the  form  of  basic  carbonate,  by  means  of  ammonium 
carbonate.  The  basic  carbonate  is  subsequently  converted  into 
plumbic  oxide,  PbO,  by  heat,  and  the  metal  is  weighed  in  this 
form.  Factor — 

(PbO)  223  :  (Pb)  207  =  i  :  0*92825 

Estimation  of  Lead  in  Lead  Acetate,  Pb(C2H3O2)2,3H2O. 
— Employ  about  0-75  gram.  The  recrystallised  salt  is  weighed  out 
into  a  beaker,  and  dissolved  in  water  with  the  addition  of  a  few  drops 
of  acetic  acid.  The  mixture  is  slightly  warmed,  and  a  solution  of 
ammonium  carbonate  containing  a  little  ammonia  is  added,  until 
precipitation  is  complete.  Excess  of  the  reagent  is  detrimental,  as 
the  precipitated  lead  carbonate  is  somewhat  soluble  in  ammoniacal 
solutions.  The  solution  should  be  allowed  to  become  cold  before 
being  filtered.  The  precipitate  is  washed  upon  the  filter  with  cold 
water  until  a  drop  of  the  filtrate  gives  no  alkaline  reaction  with 
turmeric  paper,  and  afterwards  dried  in  the  steam-oven. 

The  further  treatment  of  the  dry  precipitate  is  carried  out  by 
either  of  the  two  following  methods  : — 


Estimation  of  Lead. 


241 


(a)  The  precipitate  is  detached  from  the  filter  and  carefully 
deposited  upon  a  small  clock-glass,  and  covered  with  an  inverted 
beaker  or  funnel  for  protection.  The  filter  is  then  folded  up  and 
placed  in  a  porcelain  crucible,  which  is  heated  until  the  paper  is 
completely  burnt.  When  it  has  again  cooled,  the  ash  is  moistened 
with  a  drop  or  two  of  strong  nitric  acid  from  a  pipette,  after  which 
it  is  gently  warmed  to  expel  the  excess  of  acid.  (In  this  way  the 
reduced  lead  is  converted  into  lead  nitrate,  which  in  the  subsequent 
heating  is  decomposed  into  the  oxide.)  The  precipitate  upon  the 
clock-glass  is  then  carefully  transferred  to  the  crucible,  by  means  of 
a  small  camel's-hair  brush  or  a  feather,  in  the  manner  shown  in 
Fig.  49  ;  after  which  the  crucible  is  heated  by  means  of  a  Bunsen 
flame  to  a  dull-red  heat  for  about  ten  minutes.  It  is  then  placed  in 


FIG.  49. 

the  desiccator  to  cool,  and  weighed.  The  heating  is  repeated  until 
the  weight  is  constant,  care  being  taken  not  to  raise  the  temperature 
sufficiently  high  to  fuse  the  lead  oxide. 

(£)  The  precipitate  is  transferred  at  once  to  the  porcelain 
crucible.  The  filter  is  cut  into  about  three  pieces,  cutting  from  the 
apex  of  the  cone  to  the  upper  edge  with  a  clean  sharp  pair  of 
scissors.  Each  piece  is  then  separately  burnt  by  holding  it  with  a 
pair  of  tongs  over  the  inverted  crucible-lid,  supported  on  a  pipe- clay 
triangle,  in  the  manner  shown  in  Fig.  50.  The  ash  is  then 
moistened  upon  the  crucible-lid  with  a  drop  of  nitric  acid,  and  the 
lid  is  gently  warmed  to  expel  the  excess  of  acid.  The  crucible  and 
lid  are  together  heated  until  the  precipitate  is  completely  converted 
into  the  oxide  as  in  the  previous  description. 


242  Typical  Gravimetric  Estimations. 

Formula-weight     of    Pb(C2H3O2)2,3H2O  =  379.       Theoretical 
percentage  of  Pb  =  54'6i7. 


FIG.  50. 

II.  As  LEAD  SULPHATE. 

Epitome  of  Process. — The  lead  is  precipitated  as  lead 
sulphate,  PbSO4,  by  means  of  sulphuric  acid,  in  the  presence  of  a 
considerable  volume  of  alcohol.  The  precipitate  is  dried  and 
weighed.  Factor — 

(PbSO4)  303  :  (Pb)  207  =  i  :  0-68316 

Estimation  of  Lead  in  Lead  Acetate,  Pb(C2H302)2,3H2O. 
— Employ  from  o'5  to  075  gram.  The  salt  is  dissolved  in 
rather  less  water  than  the  usual  quantity,  in  order  to  keep  the 
solution  moderately  strong  (about  80  c.c.  will  be  found  convenient). 
Dilute  sulphuric  acid  is  added  until  no  further  precipitate  is 
produced.  A  moderate  excess  of  the  reagent  is  advantageous, 
since  lead  sulphate  is  less  soluble  in  dilute  sulphuric  acid  than 
in  water.  The  precipitation  is  made  complete  by  the  addition  of 
a  quantity  of  alcohol  (methylated  spirit  may  be  used)  equal  to  about 
twice  the  volume  of  the  aqueous  solution  (about  200  c.c.).  After 
the  precipitate  has  settled,  it  is  transferred  to  a  filter,  and  washed 
with  methylated  spirit  (in  which  lead  sulphate  is  quite  insoluble) 


Estimation  of  Zinc.  243 

until  the  filtrate  is  free  from  sulphuric  acid,  as  shown  by  the 
addition  of  barium  chloride  to  a  few  drops  of  the  liquid.  The 
precipitate  is  then  dried  in  the  steam-oven. 

The  filter  is  incinerated  according  to  either  of  the  two  methods 
described  in  the  foregoing  estimation.  The  ash,  however,  after 
being  moistened  with  nitric  acid,  is  treated  with  a  single  drop  of 
sulphuric  acid  in  order  to  convert  the  lead  nitrate  first  formed  into 
lead  sulphate.  After  the  excess  of  the  acids  has  been  evaporated, 
the  crucible  is  heated  to  dull  redness,  cooled  in  the  desiccator, 
and  weighed ;  the  heating  being  repeated  until  the  weight  is 
constant. 

Zinc. 

I.  As  ZINC  OXIDE. 

Epitome  of  Process. — The  zinc  is  precipitated  from  solutions 
of  its  salts  in  the  form  of  basic  zinc  carbonate  *  by  means  of  sodium 
carbonate.  The  zinc  carbonate  is  afterwards  converted  by  heat 
into  zinc  oxide,  ZnO,  in  which  form  the  metal  is  weighed.  Factor — 

(ZnO)8i  :  (Zn)65  =  i  :  0*80247 

Estimation  of  Zinc  in  Zinc  Sulphate,  ZnSO4,;H2O. — 
Employ  from  075  to  I  gram.  The  recrystallised  salt  is  weighed  out 
and  dissolved  in  a  beaker  f  in  the  usual  manner.  The  solution  is 
heated  to  boiling,  and  a  solution  of  sodium  carbonate  is  gradually 
added  until  precipitation  is  complete.  The  reagent  should  be  care- 
fully added,  so  as  to  avoid  the  loss  of  any  of  the  solution  by  the 
undue  rapidity  of  the  effervescence  due  to  the  escape  of  the  carbon 
dioxide.  The  mixture  is  boiled  in  the  covered  beaker  for  a  few 
minutes,  after  which  the  precipitate  is  allowed  to  settle.  It  is 
washed  with  boiling  water  by  decantation  two  or  three  times,  and 
is  then  transferred  to  the  filter.  The  washing  is  continued,  with 
boiling  water,  until  a  few  drops  of  the  filtrate  give  no  precipitate 
when  gently  warmed  with  a  little  barium  chloride. 

With  the  final  washings  the  precipitate  is  collected  together  as 
much  as  possible  into  the  apex  of  the  filter. 

The  dried  precipitate  is  transferred  from  the  filter  to  a  porcelain 
crucible  (previously  heated  and  weighed),  special  care  being  taken 
— by  rubbing  the  sides  of  the  filter  one  against  the  other — that  the 
least  possible  quantity  of  the  precipitate  is  left  adhering  to  the  paper. 
During  the  process  of  incineration,  the  zinc  oxide  is  liable  to  become 

*  For  the  composition  of  this  precipitate,  see  p.  55. 
f  See  note,  p.  223. 


244  Typical  Gravimetric  Estimations. 

reduced  to  the  metallic  state,  and  the  metal  so  formed  to  burn 
away  and  be  lost.  In  order,  as  far  as  possible,  to  prevent  this  loss 
by  facilitating  the  combustion  of  the  paper,  the  filter  (after  the 
removal  of  the  precipitate)  may  be  moistened  with  a  little  solution 
of  ammonium  nitrate,  and  again  placed  for  a  short  time  in  the  steam- 
oven  to  dry.*  It  is  then  rolled  up  in  platinum  wire  and  burnt,  the 
lamp-flame  being  only  momentarily  applied  to  it,  so  as  to  avoid 
unnecessary  contact  with  the  reducing  gases  of  the  flame.  The 
ash  is  allowed  to  fall  into  the  crucible,  which  is  then  heated  to 
redness  for  ten  minutes,  with  the  lid  upon  it,  over  a  Bunsen  flame. t 
After  cooling  in  the  desiccator  it  is  weighed,  and  the  heating 
repeated  until  the  weight  is  constant. 

Formula-weight  of  ZnSO,7H2O  =  287.  Theoretical  percentage 
of  Zn  =  22-648. 

II.  As  ZINC  SULPHIDE. 

Epitome  of  Process. — The  zinc  is  precipitated  from  its  solu- 
tions by  means  of  ammonium  sulphide  as  zinc  sulphide,  ZnS.  The 
precipitate  is  dried,  and  heated  in  a  stream  of  hydrogen  in  a  Rose's 
crucible,  and  the  metal  is  weighed  as  ZnS.  Factor — 

(ZnS)  97  :  (Zn)  65  =  i  :  o-67oi 

Estimation  of  Zinc  in  Zinc  Sulphate,  ZnS04,7H2O.— 
Take  about  075  gram.  The  salt  is  dissolved  in  warm  water  in  a 
beaker,  and  a  moderate  quantity  of  ammonium  chloride  is  added.:}: 
Colourless  ammonium  sulphide  is  then  added  until  precipitation  is 
complete.  The  mixture  is  boiled  for  a  short  time  in  order  to  cause 
it  to  coagulate,  and  therefore  to  settle  more  easily.  It  is  washed  twice 
by  decantation  with  hot  water  containing  a  little  ammonium  sulphide, 
and  then  transferred  to  the  filter,  where  the  washing  is  continued 
with  the  same  liquid  until  the  filtrate  is  free  from  sulphates.  The 
final  washing  should  collect  the  precipitate  as  much  as  possible  in  the 
apex  of  the  paper.  The  precipitate  is  dried  in  the  steam-oven,  after 
which  it  is  transferred  to  a  Rose's  crucible.  The  filter  is  incinerated, 
with  the  precautions  given  above,  and  the  ash  added  to  the  pre- 
cipitate. A  little  powdered  sulphur  (redistilled)  is  sprinkled  into 

*  If  the  precipitate  has  been  detached  cleanly  from  the  paper,  this  operation 
may  be  omitted. 

f  When  heating  a  crucible  containing  compounds  of  easily  reduced  metals, 
.the  lamp  should  be  placed  at  such  a  distance  beneath  the  crucible  that  the 
flame  does  not  envelop  it  and  lick  over  the  top  ;  otherwise  reduction  of  the 
compound  by  the  heated  gases  from  the  flame  is  liable  to  result ;  and  in  such  a 
case  as  zinc,  portions  of  the  metal  itself  would  be  volatilised,  and  therefore  lost. 

J  In  general  analyses,  if  the  solution  containing  the  zinc  salt  is  acid,  it  is  first 
rendered  slightly  alkaline  by  the  addition  of  ammonia. 


Estimation  of  Manganese.   '  245 

the  crucible,  the  lid  is  put  on,  and  a  gentle  stream  of  hydrogen 
is  passed  into  the  apparatus  (see  note,  p.  237).  The  crucible  is 
first  gently  heated,  the  temperature  being  gradually  increased  to  a 
bright-red  heat,  using  a  blowpipe  if  necessary.  It  is  then  allowed 
to  cool  with  the  gas  still  passing,  until  nearly  cold,  when  it  is 
removed  to  the  desiccator,  and  finally  weighed.  It  is  reheated  with 
a  little  more  sulphur  in  the  same  manner  until  the  weight  is  uniform. 

Manganese. 

Epitome  of  Process. — The  manganese  is  precipitated  as 
manganous  carbonate  by  means  of  sodium  carbonate.  The  car- 
bonate is  afterwards  converted  by  heat  into  the  tetroxide,  Mn3O4, 
in  which  form  the  metal  is  weighed.  Factor — 

(Mn3O4)  229  :  (sMn)  165  =  I  :  07205 

Estimation  of  Manganese  in  Potassium  Perman- 
ganate, KMnO4.  —  Employ  about  075  gram.  The  powdered 
recrystallised  salt  is  weighed  out  into  a  beaker  and  dissolved  in 
water.  A  stream  of  sulphur  dioxide  is  passed  through  the  solution 
until  the  liquid  is  perfectly  colourless  and  clear.*  The  manganese 
is  then  precipitated  and  washed  exactly  as  described  for  zinc  (p.  243). 
In  this  case,  however,  two  special  sources  of  error  arise.  First, 
the  precipitation  of  the  manganese  is  never  quite  complete,  causing, 
therefore,  a  loss ;  and,  second,  the  precipitate  always  carries  down 
with  it  a  certain  quantity  of  the  alkali,  thus  giving  rise  to  gain. 
Although  by  chance  these  two  errors  might  neutralise  each  other, 
an  exact  result  can  only  be  certainly  obtained  by  adopting  the 
special  methods  for  their  elimination  given  below. 

The  filtrate,  together  with  the  whole  of  the  wash-water,  is 
evaporated  to  dryness  in  a  platinum,  nickel,  or  porcelain  dish. 
The  evaporation  may  be  conducted  at  first  over  a  rose  burner, 
provided  the  flame  be  so  adjusted  that  the  liquid  never  actually 
boils.  But  when  the  liquid  reaches  such  a  state  of  concentration 
that  salts  begin  to  deposit,  the  process  must  be  finished  upon  a 
steam-bath.  During  the  entire  process  care  must  be  taken  that  no 
particles  of  foreign  matter  fall  into  the  dish  ;  it  is  therefore  desirable 

*  This  roundabout  method  of  obtaining  a  solution  of  manganous  sulphate 
is  preferred  to  the  direct  use  of  the  crystallised  sulphate,  for  the  reason  that  the 
latter  salt  is  not  easy  to  obtain  in  a  definite  state  of  hydration.  According  to 
the  conditions  of  crystallisation,  manganous  sulphate  forms  various  hydrated 
salts.  Thus,  if  deposited  below  —  6°,  the  crystals  contain  7H2O  ;  between  7°  and 
20°  the  salt  contains  5H2O  ;  while  between  20°  and  30°  it  contains  4H2O.  When 
manganous  sulphate  is  used,  the  latter  salt  is  the  best  to  employ.  It  is  pre- 
pared by  evaporating  the  solution  at  a  temperature  about  25°. 


246 


Typical  Gravimetric  Estimations. 


to  screen  the  vessel,  either  by  supporting  over  it  a  sheet  of  filter- 
paper  by  means  of  a  glass  tube  or  rod  bent  into  a  triangle  and 
held  by  a  clamp,  as  shown  in  Fig.  51;  or  by  supporting  a  large 
inverted  funnel  over  the  dish,  taking  care  that  the  water  which 
condenses  upon  the  funnel  shall  drip  outside  the  dish. 

The  residue  is  then  treated  with  a  little  hot  water,  which  dis- 
solves the  salt,  leaving  any 
manganese  present  in  the 
form  of  the  hydrated  oxide. 
This  is  filtered  through  a 
separate  filter,  washed  with 
hot  water  until  the  wash- 
water  is  free  from  sulphates, 
and  then  placed  in  the  steam- 
oven  to  dry. 

The  contents  of  the  two 
filters  are  next  transferred  to 
a  platinum  crucible,  and  the 
two  filter-papers  incinerated 
in  a  coil  of  platinum  wire, 
one  after  the  other,  and  the 
two  ashes  deposited  in  the 
crucible.  The  crucible  is 
then  raised  to  a  bright-red 
heat  for  about  ten  minutes 
(see  note,  p.  244) ;  then 
cooled  in  the  desiccator  and 
weighed.* 

The  residue  in  the  crucible  is  treated  with  a  little  boiling  water, 
the  particles  being  reduced  to  powder  by  means  of  a  short  piece  of 
stout  glass  rod  with  smooth  ends.  The  water  is  poured  off  through 
a  filter,  and  the  residue  washed  in  this  way  by  decantation  three  or 
four  times.  It  is  then  entirely  transferred  to  the  filter,  and  thoroughly 
washed  with  hot  water  until  the  washings  cease  to  give  a  pronounced 
yellow  coloration  to  the  flame  when  heated  upon  a  loop  of  clean 
platinum  wire.  The  residue  is  then  dried  in  the  steam-oven,  and 
afterwards  transferred  to  the  crucible  ;  the  filter  is  incinerated,  and 
the  crucible  again  strongly  heated,  allowed  to  cool  in  the  desiccator, 

*  Weighing  at  this  stage  is  not  necessary.  By  doing  so,  however,  the 
student  will  ascertain  the  extent  of  the  error  which  is  due  to  the  presence  of 
alkali  in  the  precipitate,  and  will  therefore  be  able  to  judge  of  the  value  of  the 
next  operation. 


FIG.  51 


Estimation  of  Potassium.  247 

and  weighed.     The  heating  is  repeated  until  the  weight  remains 
constant. 

In  calculating  the  result,  it  must  be  remembered  that  the  ash 
from  three  filters  has  to  be  deducted  from  the  weight. 

Formula-weight  of  KMnO4  =158.  Theoretical  percentage  of 
Mn  =  34'8i. 

Nickel. 

Epitome  of  Process. — The  nickel  is  precipitated  as  nickelous 
hydroxide,  Ni(HO)2,  by  means  of  potassium  hydroxide.  The  dried 
precipitate  is  converted  by  heat  into  nickelous  oxide,  NiO,  in 
which  form  the  metal  is  weighed.  Factor — 

(NiO)  75  :  (Ni)59  =  i  :  078666 

The  precipitation  is  conducted  as  in  the  case  of  copper  (p.  235  ; 
see  also  note,  p.  222).  The  filter  is  incinerated  in  a  platinum  coil, 
and  the  precipitate  is  strongly  heated  in  a  platinum  crucible 
(avoiding  the  entrance  of  reducing  gases  from  the  flame,  see  p.  244) 
until  the  weight  is  constant. 

Nickelous  hydroxide  retains  traces  of  the  alkali  with  some  per- 
sistence ;  therefore,  when  special  exactness  is  desired,  the  precipitate 
is  afterwards  treated  as  in  the  case  of  manganese. 

A  convenient  salt  in  which  to  make  an  estimation  of  nickel  is 
the  double  ammonium  nickel  sulphate,  (NH4)2SO4,NiSO4,6H2O. 
Formula-weight  =  395.  Theoretical  percentage  of  Ni  =  I4'93- 

Cobalt. 

Epitome  of  Process.— The  cobalt  is  precipitated  as  hydroxide, 
Co(HO)2,  by  means  of  potassium  hydroxide.  The  dried  precipitate 
is  reduced  in  a  stream  of  hydrogen  in  a  Rose's  crucible,  and 
weighed  in  the  metallic  state. 

The  precipitation  is  carried  out  as  in  the  case  of  nickel.  The 
filter  is  incinerated  in  a  porcelain  crucible  along  with  the  precipitate, 
after  which  a  gentle  stream  of  dry  hydrogen  is  delivered  through 
the  Rose's  apparatus  (p.  237),  and  the  precipitate  heated  to  a  low  red 
heat  in  the  gas  for  about  ten  minutes.  It  is  allowed  to  cool  in  the 
hydrogen  atmosphere,  and  then  weighed. 

Like  the  hydroxides  of  nickel  and  of  manganese,  cobalt  hydroxide 
retains  traces  of  alkali.  The  contents  of  the  crucible  must  there- 
fore be  treated  with  hot  water  until  the  washings  are  free  from 
alkali.  The  residue  is  again  dried,  and  once  more  heated  in  a 
stream  of  hydrogen. 

Potassium. 

I.  As  POTASSIUM  SULPHATE. 

Epitome  of  Process. — The  potassium  salt  is  converted  into 
potassium  sulphate,  K2SO4,  by  treatment  with  strong  sulphuric 


248  Typical  Gravimetric  Estimations. 

acid.  The  mixture  is  evaporated  to  dryness  and  strongly  heated, 
and  the  dry  residue  of  potassium  sulphate  weighed.  The  process 
is  only  applicable  in  the  case  of  compounds  of  potassium  with 
volatile  acids  which  will  be  completely  expelled  by  the  sulphuric 
acid,  leaving  nothing  but  the  potassium  sulphate.  Such  compounds, 
for  example,  as  the  nitrate,  chloride,  bromide,  etc.  Factor— 

(K2S04)  174  :  (2K)  78  -  i  :  0-44827 

Estimation  of  Potassium  in  Potassium  Nitrate,  KNO3, 
—Take  about  0-5  gram.  The  nitre  (purified  by  recrystallisation)  is 
weighed  out  into  a  platinum  crucible,  and  a  few  drops  of  strong 
sulphuric  acid  are  added.*  The  crucible  is  then  supported  in  an 
inclined  position  upon  the  pipe-clay  triangle,  the  lid  being  placed 

so  that  the  crucible  is 
very  nearly  but  not  en- 
tirely closed,  as  shown  in 
Fig.  52.  A  gentle  heat  is 
applied  to  the  crucible, 
the  flame  being  placed 
under  the  edge  of  the 
vessel.  After  the  nitric 
acid  has  been  expelled, 
the  excess  of  sulphuric 
acid  begins  to  volatilise, 
as  seen  by  the  escape  of 
white  fumes. 

FIG    2  As  the  operation  pro- 

ceeds, the  hydrogen  po- 
tassium sulphate,  which  is  the  first  product  of  the  action,  and  which 
is  a  moderately  easily  fusible  salt,  is  converted  into  the  normal  sul- 
phate with  elimination  of  sulphuric  acid.  The  normal  salt  being 
much  more  difficult  of  fusion,  the  contents  of  the  crucible  gradually 
become  pasty,  and  finally  solid.  During  this  stage,  care  must  be 
taken  that  no  loss  by  spirting  takes  place. 

The  completion  of  this  change  requires  prolonged  exposure  to 
a  bright-red  heat,  but  in  an  atmosphere  charged  with  ammonia 
it  takes  place  much  more  quickly.  Hence,  while  the  mass  is  heated 
to  a  dull  red,  a  few  small  fragments  of  ammonium  carbonate  are 
introduced  from  time  to  time,  the  lid  being  replaced  immediately. 
The  heat  is  continued  for  a  few  minutes  after  the  addition  of  the 

*  In  the  case  of  potassium  salts  containing  acids  which  are  expelled  with 
effervescence,  the  salt  should  be  first  moistened  with  water,  and  the  sulphuric 
acid  added  drop  by  drop  so  long  as  effervescence  takes  place,  the  cover 
being  replaced  immediately  after  each  addition  of  acid. 


Estimation  of  Potassium.  249 

ammonium  carbonate,  and  the  crucible  cooled  in  the  desiccator 
and  weighed.  It  is  afterwards  re-heated  with  a  little  more  of  the 
carbonate  and  again  weighed,  the  operation  being  repeated  until 
the  weight  is  constant. 

Formula-weight  of  KNO3  =  lor.  Theoretical  percentage  of 
K  =  38-61. 

II.  As  POTASSIUM  PLATINIC  CHLORIDE,  2KCl,PtCl4. 

Epitome  of  Process. — The  potassium  is  precipitated  in  the 
form  of  the  double  chloride  of  potassium  and  platinum  (potassium 
chloroplatinate),  and  the  precipitate  is  weighed  without  alteration 
upon  a  tared  filter. 

This  process  is  only  applicable  in  the  case  of  such  potassium 
compounds  as  are  wholly  converted  into  the  chloride  by  means  of 
hydrochloric  acid.  Factor — 

(2KCl,PtCl4)  486  :  (2K)  78  =  I  :  0-1605 

Estimation  of  Potassium  in  Potassium  Chloride,  KC1. 
— Employ  about  0-25  gram.  The  recrystallised  salt  is  weighed 
out  into  a  small  porcelain  dish,  and  dissolved  in  about  TO  c.  c.  of 
water,  to  which  two  or  three  drops  of  hydrochloric  acid  have  been 
added.  A  moderately  strong  solution  of  platinum  chloride  is  added 
in  slight  excess,*  and  the  mixture  evaporated  on  a  steam-bath  until 
it  becomes  semi- solid. 

The  crystalline  residue  is  rinsed  with  alcohol,  and,  after  being 
allowed  to  settle,  the  supernatant  liquid  is  poured  off  through  a 
previously  weighed  filter  (see  p.  194).  The  precipitate  is  washed 
by  decantation  with  alcohol  two  or  three  times,  and  finally  trans- 
ferred to  the  filter,  where  it  is  further  washed  with  alcohol  (con- 
tained in  a  small  wash-bottle)  until  the  filtrate  is  free  from  any 
tinge  of  yellow  colour.  The  filter  is  then  withdrawn  from  the 
funnel,  carefully  folded,  and  placed  between  a  pair  of  watch-glasses 
to  dry  in  the  steam-oven.  It  is  then  heated  and  weighed  until  the 
weight  is  constant.f 

Formula-weight  of  potassium  chloride,  KC1  =  74'5.  Theoretical 
percentage  of  K  =  52*349. 

*  It  is  desirable  that  the  strength  of  the  solution  of  platinum  chloride  should 
be  approximately  known,  in  which  case  a  quantity  of  it  will  be  added  which 
contains  a  weight  of  PtCl4  equal  to  about  four  times  the  weight  of  potassium 
chloride  taken  for  the  analysis.  The  solution  should  be  strong,  so  as  to  avoid 
unnecessary  loss  of  time  in  evaporation.  If  the  strength  of  the  platinum 
chloride  is  unknown,  it  is  only  possible  to  ascertain  whether  excess  has  been 
added  by  observing  whether  the  supernatant  liquid,  as  it  becomes  concentrated 
by  evaporation,  has  the  yellow  colour  of  the  reagent. 

f  For  an  alternative  method  of  treating  the  precipitate,  see  Estimation 
of  ammonium  below. 


250 


Typical  Gravimetric  Estimations. 


Ammonium. 

Epitome  of  Process. — The  ammonium  salt  is  first  boiled 
with  sodium  hydroxide,  and  the  ammonia  which  is  thus  expelled 
is  absorbed  by  dilute  hydrochloric  acid.  The  ammonium,  in  the 
ammonium  chloride  so  obtained,  is  then  precipitated  as  the  double 
chloride  of  ammonium  and  platinum,  2NH4Cl,PtQ4,  by  means  of 
platinic  chloride,  and  is  finally  weighed  in  this  form.  Factor — 

(2NH4Cl,PtCl4)  644  :  (2NH4)  36  =  i  :  0-08108 

Estimation  of  Ammonium  in  Ammonium  Sulphate, 

(NH4)2SO4. — Employ  about  0-25  gram.   The  purified  salt  is  weighed 
out  into  a  flask  (F,  Fig.  53),  and  dissolved  by  the  addition  of  a 

small  quantity  (30  to  40  c.c.) 
of  water.  The  rubber  cork, 
carrying  a  dropping -funnel 
and  leading-tube,  is  then  in- 
troduced ;  the  drawn-out  end 
of  the  funnel  tube  should 
nearly,  but  not  quite,  touch 
the  surface  of  the  liquid. 

To  the  other  end  of  the 
twice-bent  leading-tube  *  the 
empty  flask  G  is  attached  by 
means  of  a  second  rubber 
cork,  through  which  passes 
the  "  scrubber"  tube  S, which 
is  packed  with  small  broken 
glass,  previously  thoroughly 
washed.  The  leading  tube 
should  be  moderately  wide 
in  the  bore,  and  at  both  ends 
it  should  be  cut  obliquely,  as 
shown  in  the  figure,  f 

When  the  apparatus  is 
thus  arranged,  a  quantity  of  dilute  hydrochloric  acid  is  introduced 
into  the  absorption  flask  G  through  the  scrubber  tube,  in  order 
to  thoroughly  wet  the  broken  glass  with  the  acid.  While  this  is 
being  done  the  tap  of  the  funnel  must  be  open,  to  allow  the  air 

*  It  is  preferable  that  the  leading-tube  should  be  in  one  piece,  as  shown  in 
the  figure,  as  every  unnecessary  joint  is  liable  to  introduce  error  through 
leakage. 

t  This  is  done  in  a  few  minutes  by  grinding  the  tube  down  upon  a  sheet  of 
moderately  coarse  emery  cloth,  which  is  freely  wetted  with  turpentine. 


S 


FIG.  53- 


Estimation  of  Ammonia.  251 

to  pass  out  (which  it  will  do  if  the  funnel  does  not  touch  the  liquid, 
as  directed).  When  the  acid  is  in,  the  obliquely  cut  end  of  the 
leading-tube  should  just  dip  into  the  liquid,  so  as  to  about  half 
close  the  orifice.  The  tap  is  then  closed,  and  the  funnel  about 
three-parts  filled  with  a  moderately  strong  solution  of  sodium 
hydroxide  (i  part  of  soda  to  5  parts  of  water). 

The  solution  in  flask  F  is  heated  nearly  to  the  boiling-point, 
and  the  caustic  soda  allowed  to  enter  from  the  funnel  drop  by  drop, 
the  liquid  being  gently  boiled  the  whole  time.*  When  the  whole 
of  the  soda  has  been  introduced,  and  the  ammonia  has  been  all 
expelled  by  continuing  the  boiling  for  a  few  minutes,  the  tap  of  the 
funnel  is  opened  (the  end  now  being  beneath  the  surface  of  the 
liquid),  and  the  gaseous  contents  of  the  flask  are  very  gently 
drawn  through  into  the  absorption  vessel  by  means  of  a  piece  of 
rubber  tube  attached  to  the  scrubber  tube.  The  apparatus  is  then 
disconnected,  and  the  contents  of  the  absorption  flask  are  trans- 
ferred to  a  porcelain  dish.  The  end  of  the  leading-tube,  the 
scrubber,  and  the  flask  are  thoroughly  washed  with  water,  and  the 
washings  added  to  the  solution  in  the  dish. 

The  requisite  quantity  of  platinum  chloride  (see  note,  p.  249)  is 
added,  and  the  solution  evaporated  upon  a  steam-bath  until  it 
becomes  semi-solid.  The  crystalline  precipitate  of  2NH4Cl,PtCl4 
is  then  washed  with  alcohol,  and  afterwards  treated  in  one  of  the 
two  following  alternative  ways  : — 

(a)  It  is  transferred  to  a  tared   filter,  dried,  and  weighed  as 
described  for  the  similar  potassium  compound. 

(b)  The  precipitate  is  washed  with  alcohol  by  decantation,  the 
washings   being   poured   off  through   an   unweighed   filter.      The 
process  is  continued  until  the  decanted  liquid  is  free  from  any  tinge 
of  yellow  colour.      The  filter,  which  contains  small  quantities  of 
the  precipitate,  is  placed  in  the   steam-oven  to  dry.     The  main 
portion  of  the  precipitate  is  rinsed  from  the  dish  into  a  platinum 
crucible  by  means  of  a  fine  alcohol  jet,  and  the  crucible  placed  in 
the  steam-oven  to  dry. 

*  Care  must  be  taken  throughout  the  whole  of  the  boiling  process  that 
none  of  the  alkaline  liquid  is  carried  over  either  as  spray  or  by  frothing.  This 
precaution  is  specially  necessary  when  the  ammonia  which  is  absorbed  in  the 
flask  G  is  afterwards  to  be  estimated  by  volumetric  methods  (q.v.}.  If  the 
leading-tube  is  of  moderately  large  bore,  and  is  cut  off  diagonally  as  directed, 
the  spray  which  is  thrown  up  will  run  back.  Should  the  mixture  show  any 
tendency  to  froth  (which  will  not  be  the  case  with  the  particular  example  here 
chosen,  but  which  is  liable  to  happen  with  some  ammoniacal  liquids),  a  few 
fragments  of  paraffin  wax  should  be  added  before  commencing  to  boil. 

In  order  to  keep  the  absorption  flask  cool,  it  may  with  advantage  be  stood 
in  a  dish  of  water  during  the  operation. 


252  Typical  Gravimetric  Estimations. 

The  crucible  with  its  dry  contents  is  then  weighed,  after  which 
the  filter  is  incinerated  in  a  platinum  coil,  and  the  ash  added.  The 
crucible  is  then  weighed  again.  The  increase  in  weight  is  due  to 
the  ash  plus  a  small  quantity  of  metallic  platinum  derived  from 
that  portion  of  the  precipitate  which  was  carried  on  to  the  filter 
during  the  process  of  washing.  This  quantity  must  therefore  be 
added  to  the  weight  of  platinum  which  is  found  by  calculation  to 
be  present  in  the  precipitate  in  the  crucible  *  (as  explained  in  the 
case  of  silver,  p.  240). 

The  result  of  the  analysis  may  be  checked  by  converting  the 
whole  of  the  precipitate  into  metallic  platinum,  and  weighing  it. 
The  crucible  is  first  gently  heated  by  means  of  a  small  flame,  and 
the  temperature  gradually  raised  to  a  bright-red  heat,  after  which 
the  crucible  is  cooled  in  the  desiccator  and  weighed. 

Formula-weight  of  (NH4)2SO4  =  132.  Theoretical  percentage 
of  (NH4)  =  27-27. 

Tin. 

Epitome  of  Process. — The  tin  is  precipitated  as  sulphide 
from  either  "  stannous "  or  "  stannic "  solutions  by  means  of 
sulphuretted  hydrogen.  The  dried  precipitate  is  afterwards  con- 
verted by  heating  in  contact  with  air,  into  tin  dioxide,  SnO2,  in 
which  form  the  metal  is  weighed.  Factor — 

(SnO2)  150  :  (Sn)  118  =  i  :  078666 

Estimation  of  Tin  in  Stannous  Chloride,  SnCl2,2H2O. 
— Employ  about  o-6  gram.  The  purified  salt  is  weighed  into  a 
beaker,  and  dissolved  in  a  small  quantity  of  dilute  hydrochloric 
acid.  The  solution  is  diluted  by  the  addition  of  about  250  c.c. 
of  moderately  warm  water,  and  sulphuretted  hydrogen  passed 
through  until  the  liquid  is  saturated  with  gas.  The  beaker  is  then 
allowed  to  stand  for  an  hour  or  two  in  a  warm  place,  or  it  may  be 
gently  heated  for  a  short  time  to  a  temperature  not  higher  than 
can  be  comfortably  borne  by  the  hand.  The  precipitate  has  a  great 
tendency  to  pass  through  the  filter,  especially  during  the  process 
of  washing.  This  is  overcome  to  a  great  extent  by  employing  a 
double  filter,  or  by  employing  a  dilute  solution  of  ammonium 
acetate  acidified  with  a  little  acetic  acid,  with  which  to  wash  the 

*  This  method  may  be  adopted  as  an  alternative  in  the  estimation  of 
potassium.  In  that  case,  however,  it  is  very  important  that  the  least  possible 
quantity  of  the  precipitate  is  carried  on  to  the  filter  during  the  process  of 
washing,  for  the  reason  that,  unlike  the  ammonium  compound,  the  potassium 
compound,  when  heated,  leaves  a  residue  of  potassium  chloride  as  well  as 
metallic  platinum. 


Estimation  of  Arsenic.  253 

precipitate.  The  dried  precipitate  is  transferred  to  a  porcelain 
crucible,  and  the  filter  incinerated  in  a  platinum  coil.  The  ash  is 
moistened  with  a  single  drop  of  strong  nitric  acid,  as  in  the  case 
of  copper  (p.  235),  and  the  excess  of  acid  expelled  by  a  gentle  heat. 
The  crucible  is  then  gradually  heated,  first  with  its  lid  on,  and 
afterwards  with  it  removed,  until  the  sulphide  is  completely  con- 
verted into  oxide  ;  the  temperature  being  finally  raised  to  a  red 
heat  by  means  of  a  blowpipe.  The  last  traces  of  sulphur  are 
expelled  by  allowing  the  crucible  to  partially  cool,  introducing  a 
few  fragments  of  ammonium  carbonate,  and  again  strongly  heating 
until  the  weight  is  constant. 

Formula-weight  of  SnCl2)2H2O  =  225.  Theoretical  percentage 
of  Sn  =  52'44. 

Arsenic. 

I.  As  ARSENIOUS  SULPHIDE,  As2S3. 

Epitome  of  Process.— The  arsenic  is  precipitated  from  a 
solution  of  an  "  arsenious  "  compound  in  the  form  of  arsenious 
sulphide,  As.,S3,  by  means  of  sulphuretted  hydrogen.  The  pre- 
cipitate is  weighed  upon  a  tared  filter.  It  is  afterwards  subjected  to 
treatment  with  carbon  disulphide  to  remove  free  sulphur.  Factor — 

(As2S3)  246  :  (2 As)  150  =  i  :  0-60975 

Estimation  of  Arsenic  in  Arsenious  Oxide,  As4O6.— 
Employ  about  0-5  gram.  The  arsenious  oxide  (purified  by  re- 
sublimation)  is  weighed  out  into  a  conical  flask  capable  of  holding 
about  half  a  litre,  and  fitted  with  a  caoutchouc  cork  carrying  two 
tubes  for  the  transmission  of  a  stream  of  gas  through  the  solution, 
the  longer  tube  reaching  nearly  to  the  bottom  of  the  flask. 

The  arsenious  oxide  is  dissolved  in  a  small  quantity  (30  to 
40  c.c.)  of  dilute  hydrochloric  acid,  the  operation  being  aided 
by  gently  warming  the  mixture  upon  a  steam-bath.  The  solution 
must  not  be  boiled.  When  the  oxide  has  entirely  dissolved,  the 
liquid  is  largely  diluted  by  the  addition  of  about  200  c.c.  of  warm 
water,  and  a  gentle  stream  of  sulphuretted  hydrogen  passed  through, 
until  precipitation  is  complete. 

Precipitated  in  this  manner  from  a  warm  solution,  the  arsenious 
sulphide  settles  quickly,  and  is  easily  filtered.*  The  excess  of 

*  When  conditions  exist  which  cause  the  precipitation  of  sulphur,  such  as 
the  presence  of  arsenic  acid,  chromates,  ferric  salts,  etc.,  the  solution  should 
be  cold,  or  only  slightly  warm,  as  from  a  hot  solution  the  sulphur  is  deposited 
in  coagulated  particles,  and  in  this  form  it  is  much  more  difficult  to  dissolve  in 
carbon  disulphide.  In  this  case  the  filter  containing  the  dried  precipitate  must 
be  digested  with  carbon  disulphide  in  a  dish  upon  a  water-bath.  The  carbon 
disulphide  must  itself  be  free  from  dissolved  sulphur,  and  should  leave  no 
residue  when  a  small  quantity  of  it  is  evaporated  upon  a  watch-glass. 

S 


254  Typical  Gravimetric  Estimations. 

sulphuretted  hydrogen  is  expelled  from  the  solution  by  passing  a 
brisk  stream  of  carbon  dioxide  through  the  still  warm  solution, 
after  which  the  precipitate  is  transferred  to  a  weighed  filter,  and 
washed  with  warm  water  containing  a  little  sulphuretted  hydro- 
gen. If  any  of  the  precipitate  has  deposited  itself  upon  the  glass 
tube  delivering  the  sulphuretted  hydrogen,  it  may  be  dissolved  off 
by  means  of  ammonia  in  a  small  test-tube,  and  reprecipitated  by 
the  addition  of  hydrochloric  acid,  and  added  to  the  main  portion. 
The  precipitate,  after  being  washed  until  the  washings  are  free 
from  hydrochloric  acid,  is  dried  in  the  steam-oven,  and  weighed 
in  a  stoppered  tube  (p.  194). 

The  filter  is  then  carefully  replaced  in  the  funnel,  and  the 
precipitate  rinsed  with  small  quantities  of  carbon  disulphide,  until 
a  few  drops  of  the  liquid,  when  evaporated  upon  a  watch-glass, 
leave  no  residue.  It  is  then  once  more  dried  and  weighed. 

Formula-weight  of  As4O6  =  396.  Theoretical  percentage  of 
As  =  7575. 

II.  As  MAGNESIUM  PYRO-ARSENATE,  Mg2As2O7. 

Epitome  of  Process. — The  arsenic  is  precipitated  from 
solutions  of  "  arsenic "  *  compounds  in  the  form  of  ammonium 
magnesium  arsenate,  (NH4)MgAsO4,6H2O,  by  means  of  magnesium 
sulphate  in  the  presence  of  ammonia  and  ammonium  chloride. 
The  precipitate  is  converted  by  heat  into  magnesium  pyro-arsenate, 
Mg2As2O7,  in  which  form  the  arsenic  is  weighed.  Factor — 

(Mg2As2O7)  310  :  (2As)  150  =  i  :  0-4838 

Estimation  of  Arsenic  in  Hydrogen  Disodium  Arse- 
nate, HNa2AsO4,i2H2O. — Employ  about  075  gram.  The  re- 
crystallised  salt  is  weighed  out  into  a  wide-mouth  stoppered  bottle, 
about  500  c.c.  capacity,  and  dissolved  in  cold  water.  A  moderate 
quantity  of  ammonia  is  added,  and  then  "  magnesia  mixture  "  t  in 
excess.  The  mixture  is  then  briskly  shaken  in  the  stoppered  bottle 
for  about  five  minutes,  after  which  it  is  filtered,  washed,  and  dried, 
exactly  as  directed  for  the  treatment  of  magnesium  phosphate 

(P-  233)- 

The  dried  precipitate  is  detached  from  the  filter  and  deposited 
upon  a  clock-glass,  or  a  piece  of  black  glazed  paper,  and  covered 

*  "Arsenious"  compounds  must  be  first  oxidised  to  the  "arsenic"  con- 
dition by  gently  warming  the  solution  with  strong  hydrochloric  acid,  and 
adding  a  few  crystals  of  potassium  chlorate.  The  liquid  is  kept  at  a  gentle 
heat  until  it  no  longer  smells  of  chlorine,  and  then  rendered  alkaline  by  the 
addition  of  ammonia. 

•j-  See  Reagents,  in  the  Appendix. 


Estimation  of  Antimony.  255 

over  with  a  clean  beaker.  The  filter  is  then  moistened  with  a  few 
drops  of  solution  of  ammonium  nitrate,  and  returned  to  the  steam- 
oven.  As  soon  as  it  is  dry,*  it  is  incinerated  in  a  porcelain  crucible, 
the  temperature  being  gradually  raised  to  a  bright-red  heat.  The 
crucible  is  then  allowed  to  cool,  and  the  precipitate  transferred  to  it 
(see  p.  241),  after  which  it  is  heated  in  a  gentle  stream  of  oxygen  by 
means  of  Rose's  apparatus.  The  heat,  which  is  applied  very  gradu- 
ally at  first,  is  finally  increased  to  bright  redness.  After  being  cooled 
and  weighed,  it  is  reheated  in  oxygen  until  the  weight  is  constant. 
Formula-weight  of  HNa2As04,i2H2O  =  402.  Theoretical  per- 
centage of  As  =  18*65. 

Antimony* 

Epitome  of  Process. — The  antimony  is  precipitated  as 
antimonious  sulphide,  Sb.2S3,  by  means  of  sulphuretted  hydrogen. 
The  precipitate,  which  retains  water  after  drying  in  the  steam-oven, 
and  which  also  usually  contains  free  sulphur,  is  treated  by  one  of 
the  following  methods  : — 

(a)  It  is  freed  from  water  and  sulphur  by  being  heated  in  a 
porcelain  boat  in  a  stream   of  carbon  dioxide,  and   weighed  as 

r»i       r* 

ob2S3 ;  or 

(b)  It  is  converted  into  the  tetroxide,  Sb2O4,  by  treatment  with 
strong  nitric  acid,  and  weighed  in  this  form.f     Factors — 

(Sb2S3)336  :  (2Sb)  240  =  i  :  071428 
(Sb204)  304  :  (2Sb)  240  =  i  :  078947 

Estimation  of  Antimony  in  Tartar  Emetic,  (SbO)K- 
(C4H4O6).  Employ  about  075  gram. — The  dried  salt  is  weighed 
into  a  conical  flask,  and  dissolved  in  a  little  water.  Three  or  four 
grams  of  tartaric  acid,  dissolved  in  a  little  water,  are  added,  and 
then  a  few  drops  of  hydrochloric  acid.  The  solution  is  then 
diluted  by  the  addition  of  about  200  c.c.  of  slightly  warm  water 
(if  sufficient  tartaric  acid  has  been  introduced,  this  dilution  will 
not  cause  any  precipitation),  and  a  slow  stream  of  sulphuretted 
hydrogen  passed  through,  as  described  in  the  case  of  arsenic  (p.  253). 
When  the  solution  is  saturated  with  the  gas,  the  exit  and  delivery 
tubes  may  be  joined  by  means  of  a  piece  of  rubber  tube,  in  order 
to  prevent  the  entrance  of  air,  and  the  flask  is  allowed  to  stand  in 
a  warm  place  for  an  hour  or  two. 

The  excess  of  sulphuretted  hydrogen  is  then  expelled  by  passing 

*  If  left  too  long  in  the  oven,  the  ammonium  nitrate  is  expelled,  and  the 
object  for  which  it  is  added,  namely,  to  aid  the  combustion  of  the  paper,  will 
be  defeated. 

f  This  latter  method  is  preferable  in  cases  where  the  quantity  of  antimony 
is  small. 


256  Typical  Gravimetric  Estimations. 

a  brisk  stream  of  carbon  dioxide  through  the  mixture,  which  may 
at  the  same  time  be  gradually  heated  to  the  boiling-point.*  The 
precipitate  is  transferred  to  a  weighed  filter,  and  washed  as  quickly 
as  possible  (without  allowing  the  filter  to  run  dry)  with  hot  water 
containing  a  little  sulphuretted  hydrogen.  It  is  then  dried  in  the 
steam-oven  until  the  weight  is  constant. 

The  precipitate,  consisting  of  Sb2S3  plus  an  unknown  quantity 
of  water  and  free  sulphur,  is  treated  in  one  of  the  following  ways  : — 

(a)  A  certain  proportion  of  it  (as  much  as  convenient)  is  trans- 
ferred to  a  porcelain  boat  which  has  previously  been  heated  and 
weighed,  and  the  boat  with  its  charge  again  weighed.  The  boat 
is  then  pushed,  by  means  of  a  glass  rod,  into  a  wide  glass  tube 
about  25  c.c.  (10  inches)  long,  supported  in  a  horizontal  position, 
and  a  gentle  current  of  dry  carbon  dioxide  passed  over  it.  The 
boat  is  gently  heated  by  means  of  a  Bunsen  flame  in  the  stream 
of  gas,  whereby  the  water  and  the  sulphur  are  expelled  from  the 
precipitate.  The  tube  is  allowed  to  cool  with  the  gas  still  passing 
through,  after  which  the  boat  containing  its  residue  of  black 
anhydrous  Sb2S3  is  withdrawn  and  weighed.! 

From  the  loss  of  weight  the  contents  of  the  boat  have  suffered 
by  this  treatment,  the  quantity  of  pure  Sb2S3  which  was  present 
in  the  whole  of  the  original  precipitate  can  be  deduced,  and  from 
this  the  amount  of  antimony  can  be  calculated. 

(fr)  A  certain  proportion  of  the  precipitate  is  transferred  to  a 
weighed  porcelain  crucible,  and  the  crucible  with  the  precipitate 
again  weighed.!  It  is  moistened  by  the  addition  of  three  or  four 
drops  of  ordinary  strong  nitric  acid  (sp.  gr.  1*42),  and  the  crucible 
covered  with  an  inverted  lid.  By  means  of  a  pipette  3  or  4  c.c. 
of  fuming  nitric  acid  (sp.  gr.  1*5)  are  introduced,  the  lid  being 
drawn  slightly  aside  and  quickly  replaced.  Rapid  oxidation  of  the 
antimony  sulphide  and  free  sulphur  immediately  takes  place,  after 
which  the  contents  of  the  crucible  are  evaporated  to  dryness  upon 
a  steam -bath. 

The  crucible  is  then  heated  by  means  of  a  Bunsen  flame,  at 

*  Since  antimonious  sulphide  is  decomposed  by  boiling  hydrochloric  acid, 
this  operation  would  not  be  safe  if  much  acid  were  present. 

f  A  small  portion  of  the  original  dried  precipitate  may  be  tested  for  free 
sulphur  by  heating  it  in  a  test-lube  with  a  little  strong  hydrochloric  acid.  If 
it  entirely  dissolves  to  a  perfectly  clear  liquid,  no  sulphur  is  present,  in  which 
case  the  portion  which  is  treated  in  the  boat  will  only  require  to  be  very  gently 
heated  to  expel  the  water. 

J  In  cases  where  the  total  precipitate  is  very  small,  the  filter  with  its  entire 
contents  is  transferred  to  the  crucible.  In  such  a  case  the  filter-ash  must  be 
deducted  from  the  final  weight. 


Estimation  of  Cadmium.  257 

first  very  gently,  whereby  sulphuric  acid  is  expelled,  and  antimony 
tetroxide,  Sb2O4,  is  left.  The  crucible  is  cooled  in  the  desiccator, 
and  weighed  ;  the  heating  being  repeated  until  the  weight  remains 
constant.  From  this  result  the  weight  of  antimony  tetroxide  which 
would  be  yielded  by  the  whole  of  the  precipitated  sulphide  can  be 
calculated,  and  therefrom  the  weight  of  antimony. 

Formula-weight  of  (SbO)K(C4H4O6)  =  323.  Theoretical  per- 
centage of  Sb  =  37*15. 

Cadmium. 

Epitome  of  Process. — The  cadmium  is  precipitated  as  cad- 
mium sulphide,  CdS,  by  means  of  sulphuretted  hydrogen.  The 
precipitate  is  dried,  and  weighed  upon  a  tared  filter.  Factor — 

(CdS)  i44:(Cd)  112  =  1:07777 

The  precipitation  is  conducted  in  a  beaker,  from  a  solution 
acidified  with  hydrochloric  acid.  The  precipitate  is  first  washed 
with  water  containing  a  little  sulphuretted  hydrogen  and  a  few 
drops  of  hydrochloric  acid,  and  afterwards  with  plain  water. 

The  dried  precipitate,  after  weighing,  is  extracted  with  carbon 
disulphide  to  remove  traces  of  free  sulphur  (see  Arsenic,  p.  254),  and 
again  dried  and  weighed. 

Mercury. 

Epitome  of  Process. — The  mercury  is  precipitated  in 
the  form  of  mercuric  sulphide,  HgS,  by  means  of  sulphuretted 
hydrogen.  The  precipitate  is  dried,  and  weighed  upon  a  tared 
filter.  Factor — 

(HgS)  232  :  (Hg)  200  =  i  :  0-86206 

The  details  of  the  operation  are  the  same  as  in  the  case  of 
cadmium. 


SECTION    III. 
TYPICAL  GRAVIMETRIC  ESTIMATIONS  OF  ACID  RADICALS. 

Sulphuric  Acid,  (SO,). 

Epitome  of  Process. — The  sulphuric  acid  radical  is  precipi- 
tated in  the  form  of  barium  sulphate,  BaSO4,  by  means  of  barium 
chloride  in  the  presence  of  hydrochloric  acid.  The  precipitate  is 
dried,  and  weighed  in  the  same  form.*  Factor — 

(BaS04)  233  :  (S04)  96  -  i  :  0-4120 

Estimation  of  (SO4)  in  Potassium  Alum,  K2SO4,A12- 
(SO4)3,24H2O  t — Employ  about  0*5  gram.  The  salt  is  weighed  out 
into  a  beaker,  dissolved  in  water,  and  the  solution  heated  to  the 
boiling-point.  A  small  quantity  of  hydrochloric  acid  is  added, 
and  then  a  little  moderately  strong  solution  of  ammonium  chloride. 
While  the  liquid  is  still  boiling,  a  solution  of  barium  chloride, 
heated  to  boiling  in  a  large  test-tube,  is  added  until  precipitation 
is  complete.  The  precipitate  is  washed  three  or  four  times  by 
decantation  with  hot  water,  and  finally  transferred  to  the  filter, 
where  it  is  washed  until  the  nitrate  is  free  from  chlorides.  The 
last  washings  should  be  made  to  collect  the  precipitate  as  much  as 
possible  into  the  apex  of  the  filter,  after  which  it  is  placed  in  the 
steam-oven  to  dry. 

The  dry  precipitate  is  detached  as  completely  as  possible  from 
the  paper  and  transferred  to  a  platinum  crucible,  and  the  filter 
incinerated  in  a  platinum  coil.  The  precipitate  should  be  shaken 
slightly  to  one  side  in  the  crucible,  so  that  the  ash  may  be  deposited 
in  a  clear  space  upon  the  bottom.  By  means  of  a  fine  pipette,  the 
ash  is  moistened  with  a  single  drop  of  hydrochloric  acid,  in  order 
to  convert  into  barium  chloride  any  sulphide  which  may  have  been 

*  Barium  is  estimated  as  BaSO4  by  precipitation  with  dilute  sulphuric  acid, 
the  process  being  carried  out  as  described  above.  Strontium  is  also  estimated 
in  the  same  manner,  except  that  for  the  complete  precipitation  of  SrSO4  it  is 
necessary  to  add  a  quantity  of  alcohol  (methylated  spirit)  equal  in  bulk  to  the 
volume  of  the  aqueous  liquid. 

f  For  the  treatment  of  insoluble  sulphates,  see  p.  184 


Estimation  of  Carbon  Dioxide. 


259 


formed  by  the  reduction  of  a  portion  of  the  sulphate  adhering  to 
the  paper.  Then  a  drop  of  dilute  sulphuric  acid  is  similarly  added, 
and  the  crucible  very  gently  heated  to  expel  the  excess  of  the  acids, 
after  which  the  temperature  is  raised  to  a  red  heat.  It  is  then 
cooled  in  the  desiccator  and  weighed.* 

Formula-weight  of  K2SO4,A12(SO4)3,24H2O  =  948.     Theoretical 
percentage  of  SO4  =  40-50. 

Chlorine  in  Chlorides  t 

Epitome  of  Process. — The  chlorine  is  precipitated  in  the 
form  of  silver  chloride,  AgCl,  by  means  of  silver  nitrate,  and  is 
weighed  in  this  form.  The  filter  is  incinerated  apart  from  the 
precipitate  (see  Silver,  p.  238).  Factor — 

(AgCl)  142-5  :  (Cl)  35'5  -  i  :  0-2491 

Estimation  of  Chlorine  in  Sodium  Chloride,  NaCl. 
Employ  about  0*25  gram. 

The  process  is  carried  out  as  in  the  estimation  of  silver,  p.  238. 
Formula-weight   of  NaCl  =  58-5.      Theoretical   percentage   of 
Cl  =  6c-68. 

Carbon  Dioxide. 

I.  ESTIMATION  BY  DIFFERENCE. 

Epitome  of  Process. — A  known  weight  of  the  carbonate 
is  decomposed  by  a  suitable  acid  in  a 
weighed  apparatus,  and  the  carbon  dioxide 
expelled.  The  loss  of  weight  which  re- 
sults, represents  the  carbon  dioxide  in  the 
compound.  Various  forms  of  apparatus 
are  used.  Fig.  54  represents  a  simple 
apparatus  which  can  readily  be  fitted  up 
by  the  student  himself,  and  which  gives 
very  good  results.  It  consists  of  a  small 
flask  of  thin  glass  fitted  with  a  cork  carry- 
ing a  thistle  funnel  and  a  bent  exit  tube. 
The  exit  tube  is  drawn  down  to  a  mode- 
rately fine  tube,  which  passes  to  the  bottom 
of  a  small  narrow  test-tube,  /.  The  portion 

*  In  the  estimation  of  (SO.,)  as  barium  sulphate 
in  general  analysis,  whether  in  a  sulphate,  or  in  a 
sulphide  after  oxidation  into  (SO4),  the  following 
points  must  be  borne  in  mind  :  nitric  acid, 
nitrates,  and  chlorates  must  be  absent ;  hydro- 
chloric acid  must  not  be  present  in  large  quantity.  J 

f  The  estimation  of  bromine  in  bromides  is 
carried  out  in  the  same  manner,  the  precipitate 
being  treated  by  the  alternative  method,  No.  2, 
given  on  p.  239.  L,IG 


260 


Typical  Gravimetric  Estimations. 


of  the  funnel  tube  which  extends  down  into  the  flask  is  also  drawn 
out.  A  short  piece  of  small  rubber  tube  is  introduced  into  the 
thistle  funnel  at  c,  and  a  piece  of  drawn-out  glass  rod,  s,  pushed 
into  this,  serves  as  a  stopper.  Dilute  acid  (for  decomposing  the 
carbonate)  is  introduced  into  the  funnel,  and  the  tube  /  is  half 
filled  with  strong  sulphuric  acid  in  order  to  arrest  the  escape  of 
aqueous  vapour  along  with  the  issuing  carbon  dioxide.  By  slightly 
lifting  the  little  stopper  s,  the  admission 
of  acid  can  be  regulated  with  consider- 
able nicety. 

The  apparatus  shown  in  Fig.  55 
(known  as  the  Schrotter  apparatus) 
possesses  the  advantage  of  not  con- 
taining any  corks.  The  carbonate  is  in- 
troduced at  the  stoppered  neck  s  ;  the 
acid  is  admitted  from  the  stoppered 
funnel  /,  while  the  escaping  carbon 
dioxide  is  dried  by  its  passage  through 
the  strong  sulphuric  acid  contained  in  d. 
The  gas  passes  up  the  central  tube 
within  d,  as  indicated  by  the  arrows, 
and  forces  the  acid  down  to  the  level 
of  the  holes  near  the  bottom  of  the 
second  tube,  and  then  bubbles  out 
through  the  acid  and  escapes  at  the 
top  of  the  outer  tube. 

Estimation  of  CO  in  Potas- 
sium Carbonate,  K2C03.— Employ 
about  i'5  gram.  The  pure  dry  salt  is 
weighed  into  the  flask  of  the  apparatus, 
and  the  slightly  greased  stopper  in- 
serted in  the  neck  s.  The  dropping- 
funnel  is  then  nearly  filled  with  dilute 

sulphuric  acid,  and  the  drying-tube  is  about  half  filled  with  strong 
oil  of  vitriol,  and  the  apparatus  carefully  weighed.  The  acid  from 
the  funnel  is  then  allowed  to  very  slowly  enter  the  flask,  so  that  the 
carbon  dioxide  may  escape  in  single  bubbles  through  the  strong 
acid  in  the  drying-tube.  If  the  bubbles  pass  in  a  stream  faster  than 
about  two  to  the  second  (i.e.  too  quickly  to  be  easily  counted),  the 
gas  will  not  be  deprived  of  aqueous  vapour  by  the  acid,  and  the 
result  of  the  analysis  will  be  vitiated.  When  the  effervescence  has 
ceased,  the  remainder  of  the  acid  is  allowed  to  run  in  and  the  tap 
left  open.  The  flask  is  then  very  gently  heated  over  a  small  flame, 
in  order  to  expel  the  whole  of  the  carbon  dioxide,  until  the  liquid 
just  begins  to  boil.  During  the  heating,  the  gaseous  contents  of  the 


Estimation  of  Carbon  Dioxide.  261 

flask  are  slowly  sucked  out  by  means  of  an  aspirator  attached  by 
a  caoutchouc  tube  to  the  exit  of  the  drying-tube.  The  aspirator 
shown  on  p.  204  may  be  employed,  or  the  gas  may  be  drawn  out 
by  the  lungs.  In  the  latter  case,  the  peculiar  sweet  taste  of  the 
carbon  dioxide  makes  it  easy  to  tell  when  the  whole  of  the  gas  has 
been  withdrawn  and  its  place  taken  by  air.  The  aspiration  must 
not  be  unnecessarily  prolonged.  The  apparatus  is  then  allowed 
to  become  cold,  and  afterwards  weighed.  While  the  apparatus  is 
cooling,  the  exit  tube  of  the  drying-vessel  d  should  be  closed  with 
a  cap  (consisting  of  a  short  piece  of  caoutchouc  tube  plugged  at 
one  end  with  a  fragment  of  glass  rod)  in  order  to  prevent  the 
absorption  of  atmospheric  moisture  by  the  sulphuric  acid  ;  the  cap 
is  removed  while  weighing.* 

Formula-weight  of  K2CO3  =  138.  Theoretical  percentage  of 
CO,  =  31-88. 

II.  ESTIMATION  BY  DIRECT  WEIGHING. 

Epitome  of  Process- — The  carbonate  is  decomposed  by  an 
acid,  and  the  evolved  carbon  dioxide  is  absorbed  by  soda-lime  in 
a  weighed  apparatus. 

Fig.  56  represents  the  apparatus  which  may  be  employed  for 
the  process.  The  carbonate  is  decomposed  in  the  flask  B  by  means 
of  acid  introduced  from  the  dropping-funnel,  the  fine  drawn-out 
end  of  which  reaches  nearly  to  the  bottom.  The  carbon  dioxide 
is  first  dried  by  passing  through  the  tube  E,  containing  pumice 
moistened  with  sulphuric  acid.  It  is  then  deprived  of  acid  vapours 
by  means  of  the  tube  D,  filled  with  pumice  impregnated  with  an- 
hydrous copper  sulphate  (note  below).  The  pure  dry  gas  then 
passes  through  the  weighed  absorption- tube  A.  which  is  about 
four-fifths  filled  with  soda-lime,  the  remaining  portion,  a,  furthest 
removed  from  the  incoming  gas,  being  filled  up  with  dry  calcium 
chloride.  Beyond  the  absorption-tube  is  a  smaller  U-tube  contain- 
ing soda-lime,  which  is  weighed  immediately  before  and  after  the 
experiment.  This  serves  as  a  guard  tube,  and  its  weight  should 
remain  constant  if  the  absorption- tube  A  is  properly  absorbing  the 
whole  of  the  carbon  dioxide.  After  this  guard  tube  is  placed  a 

*  The  accuracy  of  the  results  obtained  by  this  method  depends  very  largely 
upon  the  slowness  with  which  the  carbon  dioxide  is  disengaged.  The  method 
is  more  suitable  for  those  carbonates  which  give  up  the  whole  of  their  carbon 
dioxide  when  acted  upon  by  sulphuric  acid  ;  in  other  words,  carbonates  of  those 
metals  which  yield  soluble  sulphates.  When  hydrochloric  or  nitric  acid  is 
employed,  an  additional  source  of  error  is  introduced  by  the  volatilisation  of 
small  quantities  of  the  acid  during  the  process  of  expelling  the  carbon  dioxide. 
This  may,  however,  in  a  measure  be  prevented  by  attaching  to  the  drying-tube 
d  another  small  thin  glass  tube  rilled  with  particles  of  pumice  which  have  been 
soaked  in  strong  copper  sulphate  solution,  and  afterwards  heated  in  an  air-oven 
to  200°.  The  anhydrous  copper  sulphate  absorbs  hydrochloric  acid  as  well 
as  aqueous  vapoiir. 


262 


Typical  Gravimetric  Estimations. 


U-tube  containing  sufficient  strong  sulphuric  acid  to  cover  the  bend. 
This  serves  the  double  purpose  of  preventing  the  entrance  of 
moisture  to  the  guard  tube,  and  also  as  an  indicator  of  the  rate  at 
which  the  gas  is  passing  through  the  tubes.  The  tube  S,  contain  - 


FIG.  56. 

ing  soda-lime,  is  for  the  purpose  of  removing  carbon  dioxide  from 
the  air  which  is  drawn  through  the  apparatus  at  the  conclusion  of 
the  operation  by  means  of  the  aspirator  (see  p.  204).  This  tube 
and  the  aspirator  are  not  connected  until  the  final  stage  of  the 
process. 

Estimation  of  CO2  in  Calcium  Carbonate,  CaCO3.— Em- 
ploy about  i  gram.  The  pure  dry  carbonate  is  weighed  out  into 
the  flask  B.  The  caoutchouc  stopper  carrying  the  funnel  and 
exit-tube  is  inserted,  and  the  funnel  rilled  with  dilute  hydrochloric 
acid.  The  apparatus  is  then  connected  up  as  shown  in  the  figure, 
the  absorption-tube  A  having  previously  been  weighed.*  The  acid 
is  allowed  slowly  to  enter  the  flask,  the  admission  of  the  acid  being 
regulated  by  the  rate  at  which  bubbles  are  seen  to  pass  through 
the  acid  in  the  last  tube  ;  they  should  not  escape  faster  than  two 

*  It  will  be  obvious  that  the  apparatus  must  be  perfectly  free  from  leakage. 
Before  being  used,  it  may  be  tested  by  connecting  a  long  vertical  glass  tube, 
dipping  into  a  vessel  of  water,  to  the  exit  of  the  last  U-tube,  and  then  sucking 
air  out  of  the  apparatus,  through  the  dropping-funnel,  until  a  column  of  water 
has  been  drawn  up  the  long  tube.  On  closing  the  tap  of  the  funnel,  the  water 
in  the  tube  should  remain  stationary.  . 


Estimation  of  Nitric  Acid.  263 

per  second.  When  effervescence  is  at  an  end,  the  remainder  of 
the  acid  is  run  in,  the  tap  is  left  open,  and  the  tube  S  is  attached 
to  the  funnel.  The  aspirator  is  connected  to  the  other  end,  and 
arranged  to  draw  a  very  slow  stream  of  air  through  the  apparatus. 
A  gentle  heat  is  applied  to  the  flask,  and  the  temperature  of  the 
liquid  kept  for  a  short  time  just  below  the  boiling-point.  The 
aspiration  should  be  continued  for  1 5  to  20  minutes,  to  ensure  the 
whole  of  the  carbon  dioxide  being  drawn  out  of  the  flask  and  ab- 
sorbed in  the  soda-lime. 

During  this  absorption  the  soda-lime  becomes  perceptibly  warm, 
and  loss  would  result  from  the  esca-pe  of  moisture  from  the  material. 
This  loss  is  prevented  by  the  layer  of  calcium  chloride  contained  in 
the  tube. 

When  the  process  is  concluded,  the  aspirator  is  stopped,  and  the 
absorption-tube  A  disconnected.  The  ends  of  the  exit-tubes  are  at 
once  closed  with  caps,  and  the  tube  carefully  wiped  and  weighed. 
The  caps  are  removed  during  the  weighing.* 

The  increase  in  weight  is  the  carbon  dioxide  expelled  from  the 
calcium  carbonate  employed. 

Formula- weight  of  CaCO3  —  100  ;  theoretical  percentage  of 
CO,  =  44- 

Nitric  Acid  (NO3). 

Epitome  of  Process.— The  nitrate  is  strongly  heated 
with  powdered  silicon  dioxide,  whereby  the  whole  of  the  N2O5 
present  in  the  nitrate  is  expelled,  and  the  oxide  of  the  metal 
remains  associated  with  the  silica,  f  The  loss  of  weight  represents 
N2O5,  from  which  the  proportion  of  NO3  present  can  be  calculated. 
Factor— 

(N2O5)  1 08  :  (2NO3)  124  =  i  :  1-148 

Estimation  of  (NO3)  in  Potassium  Nitrate,  KNO3.— 
Employ  0-5  to  075  gram.  From  2  to  3  grams  of  powdered  quartz 
are  placed  in  a  platinum  crucible  ;  it  is  then  strongly  heated,  and 
weighed  after  cooling.  The  dry  powdered  nitre  (previously  pre- 
pared by  just  melting  the  salt  in  a  porcelain  dish,  pouring  out 
the  fluid  into  a  clean  warm  porcelain  dish,  and  afterwards  powder- 
ing it)  is  introduced  into  the  crucible,  and  the  whole  weighed 
again.  The  nitrate  and  the  silica  are  intimately  mixed  by  means 

*  At  the  conclusion  of  the  determination,  the  various  purifying  tubes  may 
remain  connected  together,  the  extreme  exits  being  closed  with  caps ;  they  will 
then  be  ready  for  use  in  a  second  experiment. 

t  A  few  metallic  nitrates,  when  heated  alone,  give  up  the  whole  of  their 
N2O5,  and  leave  only  a  metallic  oxide;  but  most  nitrates,  when  thus  heated, 
leave  a  residue  of  uncertain  composition 


264  Typical  Gravimetric  Estimations. 

of  a  clean  stout  piece  of  wire  or  a  fine  glass  rod,  and  the  mixture 
heated  in  the  covered  crucible  for  about  20  minutes  to  a  red  heat.* 
It  is  then  allowed  to  cool  in  the  desiccator,  and  weighed  ;  the 
heating  being  repeated  until  no  further  loss  takes  place. 

Formula-weight  of  KNO3  --  101.  Theoretical  percentage  of 
(NO,) -61-38. 

Phosphoric  Acid  (POJ. 

Epitome  of  Process.— The  phosphoric  acid  is  precipi- 
tated in  the  form  of  ammonium  magnesium  phosphate  by  means 
of  "  magnesia  mixture  "  (as  in  the  estimation  of  arsenic,  p.  254). 
The  dried  precipitate  is  converted  by  heat  into  magnesium  pyro- 
phosphate,  in  which  form  it  is  weighed  (see  Estimation  of  mag- 
nesium, p.  233).  Factor — 

(Mg2P2O7)  222  :  (2PO4)  190  =  i  :  0-8558 

Silicic  Acid. 

Epitome  of  Process. — The  silicic  acid  is  precipitated  in  the 
gelatinous  state  by  means  of  hydrochloric  acid.  The  mixture  is 
evaporated  to  dryness,  whereby  the  precipitated  acid  is  converted 
into  the  anhydride  (silicon  dioxide),  SiO2.  The  silica  so  obtained 
is  then  filtered,  washed,  and  dried,  and  weighed  in  that  form. 

Silicates  which  are  decomposed  by  acids  are  treated  at  once 
with  hydrochloric  acid  ;  the  majority  are  converted  by  fusion  with 
alkaline  carbonates  into  silicates  of  the  alkalies,  which  are  after- 
wards precipitated  by  hydrochloric  acid  (see  p.  285). 

*  In  general  analysis,  where  there  is  any  likelihood  of  chlorides  or  sulphates 
being  present,  the  temperature  should  not  rise  above  that  at  which  the  crucible 
is  just  visibly  red  hot,  or  loss  will  result  from  the  decomposition  of  these  com- 
pounds. In  such  a  case,  the  heating  should  be  prolonged  to  at  least  half  an 
hour. 


SECTION    IV. 
EXERCISES  IN  GRAVIMETRIC  ANALYSIS. 

INTRODUCTION.      Sampling    and    Powdering.— When    the 

substance  to  be  analysed  is  a  metal  or  an  alloy  of  homogenous 
composition,  the  portion  taken  for  analysis  may  be  obtained  by 
boring  into  the  mass  with  a  steel  drill  in  a  lathe,  carefully  catching 
the  borings  upon  a  clean  sheet  of  paper.  If  this  plan  is  not  prac- 
ticable, the  metal  may  be  cut  with  a  clean  sharp  file,  and  the  filings 
used  for  the  analysis.  In  this  case,  however,  there  is  always  risk 
of  particles  of  iron  from  the  file  becoming  mixed  with  the  sample. 
These  can  be  removed  (unless  the  metal  happens  to  be  iron,  or 
other  metal  attracted  by  the  magnet)  by  means  of  a  magnet. 
When  the  bar,  ingot,  or  piece  of  metal  is  small,  it  may  be  broken 
or  cut  into  small  pieces.  In  the  case  of  metals,  it  is  not  necessary 
that  the  portion  taken  for  analysis  should  be  in  a  state  of  very  fine 
subdivision  ;  the  fragments  obtained  by  boring  or  even  by  break- 
ing or  cutting  are  sufficiently  small,  as  metals  for  the  most  part 
are  dissolved  without  much  difficulty.  Should  the  borings  or 
filings  contract  any  oil  or  grease  during  their  production,  this 
should  be  removed  by  washing  them  with  either  benzol  or  carbon 
disulphide. 

Solid  substances  which  are  capable  of  being  powdered,  such  as 
minerals,  ores,  slags,  etc.,  must  be  reduced  to  as  fine  powder  as 
possible,  and  a  portion  selected  for  analysis  which  represents  a  fair 
sample  of  the  whole.  A  number  of  the  selected  lumps  of  the  sub- 
stance are  first  broken  down  into  small  fragments  :  in  a  wedgwood 
mortar,  if  the  material  is  moderately  soft ;  in  a  steel  mortar,  if  hard  ; 
or  they  may  be  broken  up  by  wrapping  them  in  a  clean  cloth  and 
striking  them  with  a  hammer  upon  an  anvil.  These  fragments 
are  then  reduced  to  a  coarse  powder  ;  with  hard  substances,  the 
operation  being  performed  in  a  steel  percussion  mortar  (A,  Fig.  57). 
This  consists  of  three  parts  :  the  foot  or  base,  a ;  a  short  steel  tube, 
£,  which  fits  into  the  cavity  in  the  base  ;  and  a  steel  rod  or  plunger, 


266  Typical  Gravimetric  Separations. 

c,  which  will  just  pass  into  the  tube  or  cylinder.  A  larger  percussion 
mortar  is  shown  at  B,  Fig.  57,  in  which  a  massive  block  of  steel 
takes  the  place  of  the  parts  a  and  b.  The  fragments,  in  small 
quantities  at  a  time,  are  placed  in  the  tube,  and  the  plunger 
or  pestle  introduced.  This  is  then  sharply  struck  with  a  fairly 
heavy  hammer  several  times,  which  crushes  the  material  into 
coarse  powder.  When  all  the  lumps  chosen  have  been  crushed  in 
this  way,  the  coarse  powder  is  spread  upon  a  sheet  of  paper  and 
thoroughly  mixed  with  a  spatula  or  paper-knife.  A  small  portion 


FIG.  57- 

of  this  is  then  withdrawn,  and  reduced  to  the  finest  powder  possible 
in  an  agate  mortar.  In  carrying  out  this  last  operation,  small 
quantities  at  a  time,  taken  upon  the  end  of  a  spatula,  are  put  into 
the  mortar,  and  powdered  entirely  by  a  grinding  process,  and  not 
by  using  the  pestle  as  a  hammer,  which  would  be  liable  to  chip 
both  the  pestle  and  the  mortar.  The  powder  should  then  be  passed 
through  a  fine  sieve,  which  may  be  made  by  tying  a  piece  of  fine 
muslin  tightly  over  the  mouth  of  a  broken  beaker.*  The  portion 
which  is  retained  upon  the  sieve  should  be  returned  to  the  mortar 
for  further  grinding,  until  the  whole  of  it  is  reduced  to  a  sufficiently 
fine  powder  to  pass  through. 

*  The  little  piece  of  muslin  should  be  renewed  before  using  the  sieve  for 
another  mineral.  If  metal  sieves  are  employed,  the  greatest  care  must  be 
exercised  to  have  them  perfectly  clean.  Any  traces  of  a  former  operation 
should  be  beaten  and  brushed  out,  and  the  sieve  rinsed  by  sifting  a  small 
portion  of  the  powder  under  examination  through  it,  and  rejecting  the  siftings. 
On  no  account  should  a  sieve  be  wetted. 


and  Copper.  267 


Analysis  of  Silver  Coinage. 

(Alloy  of  Silver  and  Copper^) 

Epitome  of  Process. — The  metal  is  dissolved  in  nitric  acid, 
and  the  silver  precipitated  as  silver  chloride.  This  is  separated, 
and  weighed  as  such.  The  copper  is  precipitated  from  the  filtrate 
in  the  form  of  copper  hydroxide,  which  is  converted  by  heat  into 
copper  oxide,  and  weighed  in  this  form. 

The  silver  coin  (a  new  threepenny  piece)  is  carefully  weighed, 
and  then  dissolved  in  a  little  nitric  acid  in  a  small  conical  flask.* 
The  acid,  consisting  of  strong  nitric  acid,  to  which  about  half  its 
volume  of  water  has  been  added,  is  introduced  by  means  of  a  small 
funnel,  which  should  be  allowed  to  remain  in  the  mouth  of  the 
flask  during  the  whole  process,  as  it  serves  to  prevent  loss  of  the 
liquid  by  spirting.  The  solution  is  started  by  the  application  of  a 
gentle  heat,  after  which  the  action  continues  briskly  until  the  metal 
is  nearly  all  dissolved.  A  few  more  drops  of  acid  may  be  added, 
if  necessary,  for  the  complete  solution  of  the  coin,  but  excess  of 
acid  is  to  be  avoided.! 

When  the  metal  has  entirely  dissolved,  the  excess  of  acid  is 
evaporated  off  by  gently  boiling  the  solution.  The  liquid  is  then 
diluted  with  water  and  transferred  to  a  beaker,  the  flask  and  funnel 
being  thoroughly  rinsed  once  or  twice  with  distilled  water,  so  that 
the  whole  of  the  solution  is  transferred  without  loss. 

From  this  solution  the  silver  chloride  is  precipitated,  and  the 
silver  estimated  in  the  manner  described  on  p.  238.  The  filtrate, 
together  with  the  washings  containing  the  copper,  is  heated  to 
boiling,  and  the  copper  is  estimated  as  described  on  p.  235. 

A  more  expeditious  way  of  carrying  out  the  analysis  is  as  follows. 
The  metal  is  dissolved  in  a  stoppered  assay  bottle,  the  operation 
being  carried  out  exactly  as  described  above,  except  that  the  bottle 

*  English  silver  coinage  consists  of  925  parts  of  Ag  and  75  of  Cu  ;  i 
gram  of  the  alloy  therefore  contains  only  0^075  Cu.  The  weight  of  a  threepenny 
piece  is  about  1*4  gram,  which  therefore  contains  0*105  gram  Cu,  yielding 
0'J33  gram  CuO.  It  is  undesirable,  therefore,  to  employ  /mthan  this  weight  of 
metal  for  the  analysis.  If  a  larger  coin  is  used,  it  should  be  flattened  out  in  a 
rolling  machine,  and  the  thin  metal  cut  into  fragments  with  a  pair  of  scissors ; 
or,  in  the  absence  of  a  roller,  the  coin  may  be  broken  by  bending  it  backward 
and  forward,  holding  it  by  means  of  two  pairs  of  pliers.  The  requisite  quantity 
is  then  weighed  out. 

t  Old  silver  coins,  especially  the  larger  pieces,  sometimes  contain  small 
quantities  of  gold,  whicli  will  remain  as  a  black  residue  after  solution  in  nitric 
acid.  Minute  quantities  may  be  neglected  without  affecting  the  result  of  the 
analysis. 


268 


Typical  Gravimetric  Separations. 


requires  to  be  heated  with  more  care.  It  should  be  placed  upon 
a  piece  of  asbestos  cloth,  upon  a  hot  iron  plate  (p.  223),  or  may 
be  heated  on  a  steam-bath.  When  the  excess  of  acid  has  been 
evaporated,  the  liquid  is  diluted  in  the  flask,  and  the  silver  chloride 
precipitated  by  the  addition  of  hydrochloric  acid.  The  stopper  is 
then  inserted,  and  the  flask  vigorously  shaken  for  a  few  minutes. 
The  precipitate  is  thus  obtained  in  a  condition  which  enables  it  to 
settle  very  rapidly,  and  the  perfectly  clear  liquid  is  decanted  off 
into  a  beaker  as  thoroughly  as  possible.  The  flask  is  then  half 
filled  with  water,  and  the  contents  again  thoroughly  shaken,  and 
after  settling,  the  liquid  is  decanted.  After  two  such  washings  the 
whole  of  the  copper  will  have  been  washed  out,  and  the  solution 
containing  the  copper  is  boiled,  and  the  copper  estimated  as  already 
described. 

The  flask  containing  the  silver  chloride  is  then  completely  filled 


FIG.  58. 

with  water,  the  mouth  is  closed  with  the  finger,  and  the  vessel 
inverted  into  a  weighed  porcelain  crucible  filled  with  water,  in  the 
manner  shown  in  Fig.  58.  In  this  way  the  precipitate  is  transferred 
to  the  crucible,  the  whole  of  it  being  readily  caused  to  fall  down  by 
gently  tapping  the  flask.  When  the  whole  of  the  precipitate  has 
settled  to  the  bottom  of  the  crucible,  the  flask  is  first  gently  raised 
to  the  level  of  the  edge  of  the  crucible,  so  as  to  allow  one  or  two 
air-bubbles  to  enter  and  displace  a  little  water,  which  will  fill  up  and 
overflow  the  crucible,  and  then  it  is  quickly  drawn  away  sideways 


Analysis  of  Solder. 


269 


(being  held  in  the  position  shown  in  Fig.  59)  and  allowed  to 
empty.*  The  water  in  the  crucible  is  then  carefully  decanted  as 
much  as  possible,  and  the  crucible  placed  in  the  steam-oven  to 
dry.  When  dry,  the  crucible  may  be  gently  heated  until  the  silver 


FIG.  59. 

chloride  just  begins  to  melt  ;   and  after  cooling  in  the  desiccator, 
it  is  weighed. 

Usually  unglazed  crucibles  are  employed  for  this  operation.  The 
precipitate  dries  more  quickly  in  such  a  crucible,  and  the  chloride 
is  not  heated  after  it  is  removed  from  the  steam-oven. 


Analysis  of  Common  Solder. 

(Alloy  of  tin  and  lead.} 


5  precipitated  in  the  form  of  lead  sulphate, 
:id,  and  weighed  as  such.     Employ  about 


Epitome  of  Process.— The  alloy  is  acted  upon  by  strong 
nitric  acid,  which  converts  the  tin  into  stannic  oxide,  SnO2,  in  which 
form  the  tin  is  estimated.  The  lead  passes  into  solution  as  lead 
nitrate,  from  which  it  is 
PbSO4,  by  sulphuric  acic 
0*5  gram.f 

*  This  operation  is  quite  easily  accomplished  without  disturbing  the 
precipitate  in  the  crucible.  The  student  should  perform  the  experiment  of 
rilling  ah  empty  flask  with  water,  inverting  it  in  a  crucible,  and  then  withdrawing 
the  flask  as  here  directed.  He  will  find  that  there  is  no  difficulty  in  drawing 
the  flask,  aside  without  letting  an  air-bubble  pass  up,  which  would  of  course 
cause  a  disturbance  in  the  crucible,  and  throw  out  any  precipitate  it  contained. 
As  a  precaution,  the  crucible  may  be  placed  in  a  shallow  glass  basin,  so  that, 
should  any  of  the  precipitate  be  accidentally  thrown  out,  it  will  not  be  lost. 

f  Common  solder  consists  of  about  equal  parts  of  tin  and  lead,  but  as  there 
are  varieties  of  solder  in  commerce  containing  these  metals  in  different  propor- 
tions, a  duplicate  estimation  of  one  at  least  of  the  constituents  should  be  made 
as  a  check. 

T 


2/O  Typical  Gravimetric  Separations. 

Tin- — The  alloy  is  weighed  out  into  a  porcelain  dish,  and 
dissolved  in  a  small  quantity  of  moderately  strong  nitric  acid,  the 
dish  being  covered  with  a  clock-glass.  When  all  action  has  ceased, 
the  cover  is  rinsed  into  the  dish,  and  the  mixture  evaporated  nearly 
to  dryness  upon  a  steam-bath.  Water  is  added,  and  the  residue  of 
stannic  oxide  is  transferred  to  a  filter  and  washed. 

After  the  stannic  oxide  has  been  dried  in  the  steam-oven,  it  is 
transferred  to  a  porcelain  crucible,  being  detached  as  completely 
as  pos'sible  from  the  paper.  The  filter  is  incinerated  in  a  platinum 
coil,  care  being  taken  to  fold  the  soiled  portion  of  the  paper  into  the 
interior  of  the  roll.  The  ash  is  moistened  with  a  drop  of  nitric 
acid,  and  the  crucible  with  its  contents  heated,  first  gently  to  expel 
the  acid,  and  afterwards  to  a  red  heat.  It  is  then  cooled  in  the 
desiccator  and  weighed. 

Lead. — The  filtrate  containing  the  lead  is  precipitated  with 
sulphuric  acid,  and  the  lead  estimated  as  lead  sulphate  (see  p.  242). 

Analysis  of  German  Silver. 

(Alloy  of  copper,  zinc,  and  nickel,  with  traces  of  tin  and  iron.} 

Epitome  of  Process. — The  alloy  is  dissolved  in  nitric  acid, 
and  the  insoluble  residue  of  stannic  oxide  separated  and  weighed. 

The  copper  is  next  precipitated  as  sulphide  by  means  of  sul- 
phuretted hydrogen,  and  estimated  as  Cu.2S. 

The  filtrate  containing  the  zinc,  nickel  (and  iron),  is  boiled  to 
expel  sulphuretted  hydrogen,  and  a  few  drops  of  nitric  acid  added 
to  oxidise  the  iron,  which  is  then  precipitated  as  hydrated  oxide  by 
means  of  ammonia,  and  estimated  as  Fe2O3. 

The  zinc  is  next  precipitated  as  sulphide  by  means  of  hydrogen 
sodium  sulphide,  HNaS,  in  the  presence  of  potassium  cyanide,  and 
estimated  as  ZnS.* 

The  filtrate  is  acidified  with  nitric  acid  and  boiled  to  expel  the 
hydrogen  cyanide,  after  which  the  nickel  is  precipitated  as  hydroxide 
by  means  of  potassium  hydroxide,  and  estimated  as  NiO. 

Tin. — About  I  gram  of  German  silver  wire  or  borings  is  dis- 
solved in  strong  nitric  acid,  and  the  mixture  evaporated  almost  to 
dryness  upon  a  steam-bath.  The  residue  is  dissolved  in  water,  and 
the  solution  filtered  to  remove  the  traces  of  stannic  oxide.  After 
being  thoroughly  washed,  the  filter  is  neglected,  as  the  estimation 
of  the  tin  will  be  made  in  a  separate  larger  quantity  of  the  alloy.  \ 

*  For  alternative  methods  of  separating  zinc  and  nickel,  see  p.  272. 

f  In  the  case  of  alloys  in  which  tin  forms  one  of  the  intended  constituents, 
as  in  the  various  bronzes,  this  residue  should  be  treated  as  described  on  p.  272. 
Approximately,  the  composition  of  German  silver  is — Cu,  55  to  60  parts  ;  Zn,  22 
to  27  ;  and  Ni,  15  to  20.  In  many  samples  the  amount  of  tin  is  too  small  to 
be  determined.  A  qualitative  analysis  should  be  first  made,  in  order  to  ascertain 
whether  the  sample  contains  any  tin  and  iron. 


UNIVERSITY  OF  CALIFO«N 

DEPARTMENT  OF  CIVIL  ENCINEE! 
BERKELEY,  CALIFORNIA 

Analysis  of  German  Silver.  271 

Copper. — About  30  c.c.  of  strong  hydrochloric  acid  are  added 
to  the  filtrate,  and  the  solution  considerably  diluted  with  hot 
water.  It  is  then  heated  to  boiling,  and  a  stream  of  sulphuretted 
hydrogen  passed  through  for  about  20  minutes.  The  liquid  is 
then  filtered  quickly,  and  the  precipitate  washed  with  warm  water 
containing  sulphuretted  hydrogen. 

The  copper  sulphide  so  precipitated  is  liable  to  carry  down  with 
it  small  quantities  of  the  zinc,  also  as  sulphide  ;  it  is  necessary, 
therefore,  to  dissolve  it  and  re-precipitate  the  copper.  For  this 
purpose  the  precipitate  is  transferred  to  a  porcelain  dish,  and 
dissolved  in  strong  nitric  acid  previously  mixed  with  about  half 
its  volume  of  saturated  bromine  water,  and  the  solution  evaporated 
to  dryness  upon  a  steam-bath. 

The  residue  is  dissolved  in  about  30  c.c.  of  strong  hydro- 
chloric acid,  and  the  solution  rinsed  into  a  beaker  and  considerably 
diluted.  It  is  once  more  boiled,  and  a  stream  of  sulphuretted 
hydrogen  passed  through  it.  The  precipitated  copper  sulphide  is 
filtered  and  washed  as  before,  with  warm  water  containing  sul- 
phuretted hydrogen,  the  washing  being  continued  until  the  wash- 
water  is  free  from  hydrochloric  acid.  (Indicated  by  silver  nitrate, 
after  boiling  off  the  sulphuretted  hydrogen.) 

The  precipitate  is  dried,  and  treated  as  described  on  p.  237. 

Iron. — The  filtrates  and  washings  from  both  precipitations  of  the 
copper  are  mixed  together  and  briskly  boiled  in  a  covered  beaker, 
in  order  to  expel  the  sulphuretted  hydrogen,  after  which  the  liquid 
is  evaporated  until  the  volume  is  reduced  to  about  200  c.c. 

Two  or  three  drops  of  nitric  acid  are  introduced  in  order  to 
oxidise  the  iron,  after  which  a  moderate  quantity  of  ammonium 
chloride  is  added,  and  then  a  slight  excess  of  ammonia.  The  pre- 
cipitated ferric  hydroxide  is  filtered  and  washed  with  a  little  hot 
water.  It  is  then  dissolved  in  a  little  hot  hydrochloric  acid,  and 
re-precipitated  with  ammonia,  and  again  filtered  and  thoroughly 
washed.  The  precipitate  is  neglected,  as  the  iron,  if  only  present 
in  traces,  is  estimated  in  a  separate  larger  quantity  of  the  alloy. 

Zinc — The  two  filtrates  from  the  iron  precipitations  are  mixed 
and  boiled  until  the  steam  no  longer  smells  of  ammonia,  and  the 
liquid  concentrated  somewhat  by  evaporation.  A  freshly  made 
concentrated  solution  of  potassium  cyanide  is  then  added,  until 
the  yellowish  colour  produced  by  the  soluble  double  cyanide  of 
potassium  and  nickel  becomes  no  deeper  by  the  further  addition 
of  the  reagent.  Hydrogen  sodium  sulphide,  HNaS,  is  now  added, 
and  the  liquid  heated  to  boiling. 


272  Typical  Gravimetric  Separations. 

The  precipitation  of  the  zinc  sulphide,  which  begins  almost 
directly  the  liquid  is  heated,  is  complete  after  the  mixture  has 
boiled  for  a  few  minutes.  The  mixture  is  filtered,  and  the  pre- 
cipitate washed  with  hot  water,  and  estimated  as  zinc  sulphide,  as 
described  on  p.  244. 

Nickel. — The  solution  containing  the  double  cyanide  of  potas- 
sium and  nickel,  with  excess  of  sodium  sulphide,  is  evaporated 
clown  to  a  moderately  small  bulk,*  and  strong  nitric  acid  cautiously 
added  drop  by  drop  until  there  is  no  further  effervescence.  Both 
the  cyanides  and  the  sodium  sulphide  are  thereby  decomposed  ; 
the  former  with  evolution  of  hydrocyanic  acid  (hence  the  operation 
should  be  conducted  in  a  draught-cupboard),  while  from  the  latter 
sulphur  is  precipitated,  which  is  at  once  attacked  by  the  acid,  with 
the  evolution  of  brown  fumes.  The  mixture  is  boiled  until  the 
precipitated  sulphur  is  entirely  oxidised  to  sulphuric  acid,  after 
which  potassium  hydroxide  is  added  until  the  precipitation  of  the 
hydrated  oxide  of  nickel  is  complete. 

The  precipitate  is  then  treated  as  described  on  p.  247. 

Separate  Estimation  of  Tin  and  Iron. — For  this  purpose, 
about  5  grams  of  the  alloy  are  weighed  out  and  dissolved  in  nitric 
acid,  and  the  solution  evaporated  and  treated  as  described  above. 
The  residue  of  stannic  oxide  is  filtered  off,  and  estimated  as  on 
p.  270. 

The  filtrate  is  treated  as  described  above  for  the  separation  of 
copper  ;  in  this  case,  however,  one  precipitation  only  is  necessary. 
The  precipitate,  after  being  thoroughly  washed,  is  neglected.  The 
filtrate  and  washings  are  together  boiled  to  expel  the  sulphuretted 
hydrogen,  and  the  iron  precipitated  as  hydrated  oxide  by  means 
of  ammonia.  The  precipitate  is  treated  as  described  on  p.  230. 

Alternative  methods  for  the  Separation  of  Zinc  and 
Nickel,  (i)  Precipitation  of  Zinc  Sulphide  from  Acetic  Acid  solu- 
tion.— After  the  removal  of  the  tin,  copper,  and  iron  as  described 
above,  the  liquid  containing  the  zinc  and  nickel  is  acidified  with 
hydrochloric  acid,  and  evaporated  down  until  its  volume  is  reduced 
to  about  150  c.c.  Sodium  carbonate  solution  is  then  added 
until  a  slight  precipitate  persists,  after  which  acetic  acid  is  added 
until  the  precipitate  is  redissolved.f  Sulphuretted  hydrogen  is  then 

*  If  the  precipitation  of  the  zinc  has  been  made  in  too  dilute  a  solution,  a 
further  slight  separation  of  ZnS  may  take  place  during  this  concentration.  In 
that  case  the  liquid  must  be  filtered,  and  the  precipitate  either  added  to  the 
main  portion,  or  else  separately  weighed  and  the  result  added. 

f  Or  the  precipitate  may  be  redissolved  by  adding  hydrochloric  acid  drop  by 
drop,  and  then  removing  the  free  acid  by  the  addition  of  a  few  drops  of  a 


Separation  of  Zinc  and  Nickel.  273 

passed  through  the  mixture,  when  the  white  zinc  sulphide  is  alone 
precipitated.  When  precipitation  is  complete,  the  liquid  may  be 
gently  warmed  and  filtered.  The  precipitate  is  then  treated  as 
already  described.  The  filtrate  is  heated  to  expel  sulphuretted 
hydrogen,  and  the  nickel  precipitated  as  hydrated  oxide  by  means 
of  potassium  hydroxide,  and  the  process  carried  on  as  above. 

(2)  Precipitation  of  Zinc  Sulphide  from  a  Sucdnic  Acid  solution. 
— In  order  to  carry  out  the  separation  by  this  method,  the  iron 
must  previously  have  been  removed,  not  by  means  of  ammonia,  but 
with  sodium  acetate.* 

The  solution  containing  the  iron,  zinc,  and  nickel  (after  separa- 
tion of  the  copper  as  described  on  p.  271)  is  boiled  to  expel 
sulphuretted  hydrogen,  and  oxidised  with  nitric  acid  as  before.  It 
is  then  made  nearly  neutral  with  sodium  carbonate  ;  sodium  acetate 
is  added,  and  the  liquid  boiled  until  the  iron  is  completely  precipi- 
tated as  basic  acetate  (see  footnote,  p.  70).  The  precipitate  is 
filtered  off  and  thoroughly  washed,  after  which  it  may  be  neglected.f 

The  filtrate  and  washings  are  together  boiled  with  hydrochloric 
acid  until  the  acetates  are  entirely  destroyed  and  the  steam  no 
longer  smells  of  acetic  acid,  after  which  the  zinc  and  nickel  are 
together  precipitated  as  carbonates  by  the  addition  of  sodium 
carbonate.  The  liquid  is  filtered  and  the  precipitate  washed,  after 
which  it  is  dissolved  upon  the  filter  by  slowly  pouring  over  it  a  hot 
strong  solution  of  succinic  acid.  The  funnel  should  be  kept  covered 
with  a  clock-glass  to  avoid  loss  due  to  effervescence.  When  the 
precipitate  is  completely  dissolved,  and  the  filter  thoroughly  washed 
with  hot  water,  sulphuretted  hydrogen  is  passed  through  the  liquid 
while  still  hot,  until  the  zinc  is  completely  precipitated.  The  zinc 
sulphide  thrown  down  under  these  circumstances  settles  very  readily. 
The  sulphide  is  filtered  off,  and  treated  as  already  described. 

The  solution  is  boiled  to  expel  sulphuretted  hydrogen,  and  the 
nickel  precipitated  as  in  the  former  methods  by  means  of  potassium 
hydroxide. 

solution  ot  ammonium  acetate,  which,  by  double  decomposition,  forms 
ammonium  chloride  and  acetic  acid.  More  than  quite  a  small  quantity  of  the 
acetate  must  be  carefully  avoided,  or  else  the  nickel  will  be  partially  pre- 
cipitated by  the  sulphuretted  hydrogen. 

*  The  reason  for  this  is,  that  the  precipitation  of  zinc  as  carbonate,  by 
means  of  sodium  carbonate,  is  not  complete  in  the  presence  of  ammoniacal 
salts. 

f  In  cases  where  iron  is  present  in  sufficient  quantity  to  be  estimated  at 
this  stage,  this  precipitate  of  basic  acetate  must  be  dissolved  in  hydrochloric 
acid  ;  and,  after  boiling,  the  iron  is  precipitated  as  hydrated  oxide  by  means 
of  ammonia,  and  treated  as  already  described. 


274  Typical  Gravimetric  Separations. 

Analysis  of  Bronze  Coinage. 

(Alloy  of  copper,  tin,  and  zinc,  with  traces  of  lead.) 

Epitome  of  Process. — The  alloy  is  dissolved  in  nitric  acid, 
and  the  solution  evaporated  to  dryness.  The  residue  is  extracted 
with  water,  and  the  insoluble  stannic  oxide  dried  and  weighed. 

Sulphuric  acid  is  added  to  the  filtrate,  which  is  then  evaporated 
to  a  small  bulk,  and  the  lead  sulphate  separated  and  weighed. 

The  solution  is  diluted  with  water  to  a  definite  volume,  and  a 
small  measured  proportion  of  it  taken  in  which  to  estimate  the 
copper  ;  the  copper  in  this  separate  portion  being  precipitated  as 
cuprous  thiocyanate. 

In  the  main  portion  of  the  solution,  the  copper  is  separated  from 
the  zinc  by  means  of  sulphuretted  hydrogen  (the  precipitated 
copper  sulphide  being  neglected),  and  the  zinc  is  estimated  in  the 
nitrate. 

Tin. — A  clean  new  halfpenny  *  is  weighed,  and  dissolved  in 
strong  nitric  acid  in  a  covered  beaker.  The  solution  is  then  trans- 
ferred to  a  porcelain  dish,  and  evaporated  to  dryness  on  a  steam- 
bath.  The  residue  is  treated  with  water,  and  the  insoluble  stannic 
oxide  filtered  and  thoroughly  washed.  It  is  dried,  and  treated  as 
described  on  p.  270. 

The  stannic  oxide,  however,  is  liable  to  retain  traces  of  the 
copper  ;  it  should  therefore  be  submitted  to  the  following  treatment 
after  it  has  been  weighed  :  A  small  quantity  of  sulphur  (about 
equal  to  the  weight  of  the  stannic  oxide)  and  the  same  amount  of 
sodium  carbonate  are  mixed  together  in  the  crucible,  which  with 
its  cover  is  then  gently  heated  until  the  mass  fuses  and  the  excess 
of  sulphur  is  expelled.  When  cold,  the  mass  is  boiled  with  water, 
which  dissolves  the  sodium  thiostannate,  and  if  any  copper  were 
present,  it  remains  as  a  black  residue  of  copper  sulphide.  It  is 
filtered,  thoroughly  washed,  and  dissolved  in  nitric  acid.  The 
copper  is  then  precipitated  with  potassium  hydroxide,  and  estimated 
as  CuO,  the  weight  being  deducted  from  the  weight  of  the  stannic 
oxide. 

Lead.t — The  filtrate  from  the  stannic  oxide  is  acidulated  by 
the  addition  of  about  10  c.c.  of  strong  sulphuric  acid  (free  from 

*  English  bronze  coinage  consists  roughly  of  Cu  95,  tin  4,  zinc  i.  A  half- 
penny weighs  about  5^  grams.  This  weight,  while  affording  enough  material 
in  which  to  estimate  the  zinc,  is  much  more  than  is  necessary  for  the  copper 
determination,  hence  the  latter  is  estimated  in  a  small  portion  of  the  solution. 

f  It  is  seldom  that  bronze  coinage  contains  more  than  from  0^05  to  o'i  per 
cent,  of  lead.  In  5  grams  of  the  alloy,  o'i  per  cent,  of  lead  would  represent 
o'oo5  gram  of  the  metal,  or  0-007  gram  of  PbSO4,  a  quantity  far  too  small  to  be 
estimated,  and  scarcely  a  perceptible  turbidity  will  be  seen.  This  step  in  the 
analysis  may  therefore  be  omitted  without  any  appreciable  influence  upon  the 
result.  It  is  given  in  the  description  because  other  varieties  of  bronze  contain 
appreciable  quantities  of  lead,  which  would  be  separated  as  here  explained. 


Analysis  of  Bronze  Coinage.  27$ 

lead),*  and  evaporated  until  the  liquid  gives  off  vapours  of  sul- 
phuric acid  freely.  It  is  then  allowed  to  cool,  diluted  with  water, 
and  filtered.  The  precipitated  lead  sulphate  is  washed  with  dilute 
sulphuric  acid  (in  which  it  is  far  less  soluble  than  in  water)  until 
the  washings  are  free  from  copper,  when  the  vessel  containing  the 
nitrate  is  removed,  and  the  filter  is  washed  free  from  sulphuric 
acid  by  means  of  methylated  spirit.  (If  the  acid  were  not  thus 
washed  out,  the  paper  would  become  charred,  and  impossible  to 
manipulate,  when  dried  in  the  oven.  The  alcoholic  washings  are 
thrown  away.) 

The  lead  sulphate  is  dried,  and  treated  as  described  on  p.  242. 

Copper.— The  filtrate  containing  the  copper  and  zinc  is  diluted 
with  water  in  a  half-litre  flask  until  its  volume  is  500  c.c.  By 
means  of  a  graduated  pipette,  50  c.c.  of  this  solution  are  with- 
drawn and  transferred  to  a  beaker,  and  the  copper  precipitated  as 
cuprous  thiocyanate  by  means  of  ammonium  thiocyanate  in  the 
manner  described  on  p.  237. 

The  amount  of  copper  present  in  this  50  c.c.  of  the  solution 
will  obviously  be  one-tenth  of  that  present  in  the  entire  volume, 
and  if  any  copper  was  recovered  from  the  stannic  oxide  in  the  first 
separation,  this  must  be  added  on  in  order  to  obtain  the  total  weight 
of  copper  in  the  quantity  of  alloy  taken  for  the  analysis. 

Zinc. — The  main  portion  of  the  solution  (450  c.c.)  containing 
the  copper  and  zinc,  is  acidified  by  adding  about  50  c.c.  strong  hydro- 
chloric acid,  and  diluted  by  the  addition  of  300  c.c.  of  hot  water. 
The  mixture  is  heated  to  boiling,  and  the  copper  completely  pre- 
cipitated by  means  of  sulphuretted  hydrogen. 

The  precipitated  copper  sulphide  is  treated  as  described  on 
p.  271,  the  double  precipitation  being  even  more  necessary  in  this 
example,  where  the  proportion  of  copper  to  that  of  zinc  is  so  large. 
The  final  precipitate  of  copper  sulphide  is  here  neglected,  the  copper 
having  been  already  estimated. 

The  solutions  containing  the  zinc  are  then  evaporated  down  to 
moderate  bulk,  and  the  zinc  estimated  as  on  p.  244. 

The  weight  of  the  zinc  so  obtained  will  be  nine-tenths  of  that 
present  in  the  amount  of  alloy  employed  for  the  analysis. 

*  Ordinary  sulphuric  acid  contains  considerable  quantities  of  lead.  If  pure 
acid  is  not  available,  acid  must  be  used  which  has  been  previously  largely 
diluted  and  allowed  to  settle,  and  a  proportionately  larger  volume  will  be 
necessary. 


276  Typical  Gravimetric  Separations. 

Analysis  of  Dolomite. 

{Magnesium   and  calcium    carbonate,   containing    usually    small 
quantities  of  iron,  aluminium,  and  insoluble  silicious  matter]. * 

Epitome  of  Process. — The  mineral  is  dissolved  in  hydro- 
chloric acid,  and  the  solution  evaporated  to  dryness.  The  residue 
is  treated  with  hydrochloric  acid  and  water,  and  the  insoluble  por- 
tion, consisting  of  silica  and  silicates,  is  filtered  off,  washed,  and  dried. 

The  iron  and  alumina  are  precipitated  together  as  hydroxides,  by 
the  addition  of  ammonium  chloride  and  ammonia  to  the  filtrate.  The 
precipitate  is  dried,  and  weighed  as  oxides  of  iron  and  aluminium. 

The  nitrate  is  boiled,  and  the  calcium  precipitated  as  oxalate 
by  means  of  ammonium  oxalate,  the  precipitation  being  repeated 
to  ensure  complete  separation  of  the  magnesium.  The  precipitate 
is  dried  and  converted  by  heat  into  calcium  oxide,  and  weighed. 

The  filtrate  from  the  calcium  is  evaporated  to  dryness,  and 
heated  to  expel  excess  of  ammonium  salts.  The  residue  is  dis- 
solved in  hydrochloric  acid,  diluted  with  water,  and  the  magnesium 
precipitated  as  ammonium  magnesium  phosphate  by  the  addition 
of  ammonia  and  sodium  phosphate. 

The  carbon  dioxide  is  estimated  in  a  separate  portion  of  the 
mineral  by  the  method  described  on  p.  259. 

Silicious  Matter. — About  2  grams  of  the  finely  powdered 
mineral  (which  has  been  previously  dried  in  an  air-oven  at  a 
temperature  pf  about  200°  t)  is  weighed  out  into  a  porcelain  dish 
covered  with  a  clock-glass,  and  dissolved  by  the  cautious  addition 
of  dilute  hydrochloric  acid.  As  soon  as  action  ceases,  the  cover  is 
rinsed  into  the  dish,  a  few  drops  of  strong  nitric  acid  added,  and 
the  mixture  evaporated  to  complete  dryness  upon  a  steam-bath. 
The  residue  is  then  treated  with  a  little  strong  hydrochloric  acid, 
and  gently  warmed.  Water  is  then  added,  and  the  insoluble  residue 
of  silica  and  insoluble  silicates  is  transferred  to  a  filter,  where  it  is 
washed  free  from  hydrochloric  acid.  It  is  then  dried  in  the 
steam-oven. 

The  dried  filter  with  the  residue  is  folded  up  and  placed  in  a 
platinum  crucible,  where  it  is  strongly  heated  until  the  paper  is  com- 
pletely incinerated.!  The  crucible,  after  cooling  in  the  desiccator, 
is  weighed. 

*  Occasionally  manganese  is  met  with  in  varieties  of  dolomite,  but  usually 
only  in  traces,  which  may  be  neglected. 

f  The  amount  of  moisture  present  in  dolomite  is  usually  very  small.  It 
desired,  it  may  be  estimated  by  heating  about  10  grams  of  the  powder  to  200° 
in  the  air-oven. 

%  Usually  the  amount  of  silicious  matter  does  not  exceed  about  two  per 
cent.  In  cases  where  there  is  much  insoluble  residue,  it  should  be  detached 
from  the  paper,  and  the  latter  incinerated  apart  in  a  platinum  coil. 


Analysis  of  Dolomite.  277 

Aluminium  and  Iron.* — A  considerable  quantity  of  ammo- 
nium chloride  is  added  to  the  filtrate  (see  note  on  p.  224),  and  the 
solution  gently  warmed ;  ammonia  is  then  added,  the  least  pos- 
sible excess  being  used,  and  the  mixture  boiled.  The  precipitated 
hydroxides  of  aluminium  and  iron  are  then  filtered  and  washed  ; 
the  precipitate  is  dissolved  upon  the  filter  by  pouring  a  little  warm 
dilute  hydrochloric  acid  upon  it  (the  liquid  being  received  in  a 
separate  beaker),  and  the  acid  solution  thoroughly  washed  out  of 
the  paper. 

The  aluminium  and  iron  are  then  re-precipitated  from  this  solu- 
tion by  the  addition  of  a  slight  excess  of  ammonia  ;  the  precipitate 
is  filtered  and  washed  until  the  washings  are  free  from  chlorides, 
and  then  dried  in  the  steam-oven.  The  dry  precipitate,  if  small  in 
amount,  is  heated  with  the  paper  in  a  platinum  crucible ;  if  the 
amount  is  considerable,  it  is  detached  from  the  filter,  and  the  latter 
is  separately  incinerated  in  a  platinum  coil.  The  precipitate,  after 
being  heated,  consists  of  a  mixture  of  A12O3  and  Fe2O3 ;  it  is 
weighed  and  calculated  as  such.f 

Calcium. — The  mixed  filtrates  and  washings  from  the  double 
precipitation  of  iron  and  aluminium  are  heated  to  boiling,  and  the 
calcium  precipitated  by  means  of  ammonium  oxalate,  as  described 
on  p.  231.  Excess  of  the  reagent  must  in  this  case  be  used,  since, 
in  the  presence  of  magnesium  chloride,  the  precipitation  of  cal- 
cium oxalate  is  not  complete  unless  excess  of  ammonium  oxalate  is 
present.  The  reagent  may  be  added  in  a  solid  state  (instead  of  a 
hot  strong  solution,  as  directed  on  p.  231),  three  or  four  grams  of 
the  powdered  salt  being  introduced  into  the  boiling  liquid,  and  the 
mixture  stirred  for  a  few  minutes. 

The  precipitated  calcium  oxalate  carries  down  with  it  small 
quantities  of  magnesium  oxalate  ;  it  is  therefore  necessary  to  repeat 
the  precipitation.  For  this  purpose  the  precipitate  is  allowed  to 
settle,  and  the  clear  liquid  decanted  off  through  a  filter  without 
disturbing  the  precipitate  ;  the  latter  is  twice  washed  by  decanta- 
tion,  each  time  being  allowed  to  thoroughly  settle,  so  that  the  least 
possible  trace  of  it  only  is  transferred  to  the  filter.  It  is  then 
dissolved  in  the  beaker  by  the  addition  of  a  little  warm  dilute 
hydrochloric  acid,  and  reprecipitated  by  the  addition  of  ammonia 

*  Should  the  ore  contain  more  than  very  small  quantities  of  manganese 
(which,  however,  is  not  usually  the  case),  the  iron  and  aluminium  must  be 
separated  as  basic  acetates,  as  described  for  iron  on  p.  282. 

t  If  the  quantity  of  aluminium  and  iron  is  considerable,  and  it  is  desired 
to  determine  the  metals  separately,  the  precipitate  may  be  treated  as  described 
on  p.  278. 


278  Typical  Gravimetric  Separations. 

in  slight  excess,  and  a  few  drops  of  ammonium  oxalate  solution.* 
This  precipitate  is  allowed  to  settle,  and  transferred  to  the  same 
filter  employed  in  the  first  operation  ;  it  is  washed  and  dried,  and 
treated  as  described  on  p.  233. 

Magnesium, — The  filtrates  and  washings  from  the  calcium 
precipitate  are  evaporated  to  dryness  in  a  porcelain  dish.  The 
evaporation  may  be  conducted  at  first  over  a  gas-burner,  provided 
the  liquid  is  not  permitted  to  boil ;  but  as  the  liquid  begins  to 
deposit  salts,  the  further  heating  must  be  done  upon  a  steam- 
bath.f  The  dish  with  the  dry  residue  is  heated  over  a  ring 
gas-burner,  at  first  very  cautiously,  until  the  ammonium  salts  are 
expelled. 

When  the  dish  is  cold,  the  residue  is  dissolved  by  the  addition 
of  a  little  strong  hydrochloric  acid,  and  gently  warming  the  mix- 
ture. The  solution  is  diluted  with  water,  and  if  not  perfectly  clear 
it  is  filtered,  the  residue  in  this  case  being  neglected. 

Ammonia  in  slight  excess  is  added  (no  precipitate  should  form, 
as  the  ammonium  chloride  produced  by  the  neutralisation  of  the 
solution  will  be  sufficient  to  prevent  the  precipitation  of  the 
magnesium  hydroxide),  and  the  magnesium  then  precipitated  as 
ammonium  magnesium  phosphate  by  the  addition  of  hydrogen 
disodium  phosphate.  The  precipitation  and  the  treatment  of  the 
precipitate  are  described  on  p.  233. 

NOTE. — Iron  and  aluminium  are  associated  together  in  varying 
proportions  in  a  great  number  of  common  minerals  which  come 
under  analysis,  such  as  iron  ores,  clays,  felspar,  and  other  sili- 
cates, as  well  as  slags  and  other  artificial  products.  They  may  be 
separated  and  determined  by  the  following  methods  : — 

Separation  of  Iron  and  Aluminium.— (i)  The  precipi- 
tated hydroxides  are  dissolved  upon  the  filter  in  a  little  warm 
hydrochloric  acid,  and  the  solution  poured  into  a  strong  solution 
of  sodium  or  potassium  hydroxide  (free  from  alumina)  contained 
in  a  dish  (preferably  of  either  silver,  platinum,  or  nickel,  although 
porcelain  may  be  used  in  the  absence  of  these).  The  mixture  is 
then  boiled  for  two  or  three  minutes.  The  iron  is  precipitated 
as  ferric  hydroxide,  while  the  aluminium  passes  into  solution  as 
sodium  aluminate.  The  precipitate  is  filtered  off,  washed  with 

*  Or  the  second  precipitation  may  be  conducted  as  in  the  case  of  the  iron 
and  aluminium  ;  i.e.  the  precipitate  may  be  washed  upon  the  filter,  and  dis- 
solved by  the  addition  of  acid  to  it  while  in  the  funnel. 

t  Solutions  of  ammoniacal  salts  are  very  apt,  during  evaporation,  to  be 
drawn  up  through  the  crust  of  deposit  which  forms  round  the  edge  of  the 
liquid,  so  that  the  salt  is  continually  deposited  higher  and  higher  upon  the  sides 
of  the  dish,  until  eventually  it  creeps  over  the  edge.  To  prevent  this,  the 
finger,  slightly  greased  with  vaseline,  should  be  drawn  round  the  inside  edge 
of  the  dish  at  the  beginning  of  the  operation. 


Analysis  of  Dolomite.  279 

water,  then  redissolved  in  hydrochloric  acid  and  reprecipitated 
with  ammonia.*  The  ferric  hydroxide  is  again  filtered,  washed 
and  dried  in  the  usual  way,  and  weighed  in  the  form  of  sesqui- 
oxide  (p.  230).  The  two  nitrates  are  together  acidified  with  strong 
hydrochloric  acid,  and  the  aluminium  precipitated  as  hydroxide 
by  the  addition  of  ammonia  in  slight  excess.  The  precipitate  is 
washed  and  dried,  and  weighed  in  the  form  of  sesquioxide,  as 
described  on  p.  225. 

(2)  When  pure  caustic  soda  or  potash  (i.e.  free  from  alumina) 
is  not  at  hand,  the  process  may  be  modified  in  the  following  way, 
and   the   aluminium   estimated   by  difference.      The   precipitated 
hydroxides  of  the  two  metals  are  dried  and  weighed  together  as 
sesquioxides.     A  quantity  of  powdered  sodium  peroxide  f  is  then 
mixed  with   the   precipitate   in  the   crucible  (platinum),  and   the 
mixture  fused.     The  fused  mass  when  cold  is  cautiously  treated 
with  water  in  a  beaker  (as  the  aluminium  is  not  to  be  directly 
determined,  glass  vessels  may  be  used),  when  the  sodium  aluminate 
will  dissolve,  leaving  the  ferric  oxide.     This  is  then  filtered  and 
washed,  the  filtrate  being  rejected.     The  precipitate  is  then  dis- 
solved  in   hydrochloric   acid,  reprecipitated   with   ammonia,  and 
determined  in  the  usual  way.     The  weight  of  the  ferric  oxide  thus 
obtained,  deducted  from  the  weight  of  the  two  oxides  together, 
gives  the  weight  of  the  aluminium  sesquioxide. 

(3)  The  amount  of  iron  may  be  determined,  and  the  aluminium 
estimated   by  difference,  by  the  use  of  volumetric   methods  for 
estimating  iron.     The  mixed  oxides  are  weighed  as  in  the  above 
process  ;   they  are  then  dissolved  in  hydrochloric  acid,  and  the 
iron  determined  volumetrically,J  without  separating  the  aluminium. 
The  amount  of  sesquioxide  of  iron  corresponding  to  the  iron  thus 
determined,  when  deducted  from  the  original  weight  of  the  mixed 
oxides,  gives  the  weight  of  the  sesquioxide  of  aluminium. 

Analysis  of  Zinc-blende. 

(Zinc  sulphide,  liable  to  contain  zinc  carbonate  ;  and  lead, 
copper,  cadmium,  iron,  and  manganese,  as  sulphides  or  in  other 
forms.  A  previous  qualitative  analysis  must  be  made.} 

Epitome  of  Process.  —  The  ore  is  decomposed  by  strong 
hydrochloric  and  nitric  acids.  Any  lead  present  is  converted  into 
sulphate  and  weighed  along  with  the  silicious  matter,  or  gangue, 
from  which  it  is  afterwards  separated  by  solution  in  ammonium 
acetate. 

Copper  and  cadmium,  if  present,  are  precipitated  together  from 
the  solution  as  sulphides,  which  are  afterwards  separated  by 
dissolving  the  latter  in  dilute  sulphuric  acid. 

*  This  second  precipitation  is  necessary  in  order  to  remove  the  potash 
which  is  retained  by  the  ferric  hydroxide. 

t  Caustic  soda  may  be  substituted,  although  the  peroxide  is  preferable. 
See  Volumetric  methods. 


280  Typical  Gravimetric  Separations. 

Iron  is  next  separated  in  the  form  of  the  basic  acetate,*  the 
sulphates  present  having  been  first  converted  into  chlorides  by 
means  of  barium  chloride. 

The  manganese  is  precipitated  as  hydrated  peroxide  by  means 
of  chlorine  or  bromine  ;  and,  lastly,  the  zinc  is  determined  by 
precipitation  as  zinc  sulphide. 

The  sulphur  in  the  ore  is  estimated  in  a  separate  portion  by 
oxidation  into  sulphuric  acid,  and  subsequent  precipitation  as 
barium  sulphate. 

If  the  ore  contains  any  calamine  (zinc  carbonate),  the  carbon 
dioxide  is  determined  in  a  separate  portion  by  the  method  described 
on  p.  259. 

Sulphur. — About  I  gram  of  the  ore  (which  has  been  reduced 
to  as  fine  a  powder  as  possible,  and  dried  in  a  steam-oven)  is  weighed 
into  a  platinum  crucible,  and  then  mixed  thoroughly  with  five  or  six 
times  its  weight  of  powdered  sodium  peroxide,  f  the  mixing  being 
done  by  means  of  a  thin  glass  rod  or  stout  platinum  wire.  The 
covered  crucible  is  then  cautiously  heated  by  means  of  a  small  flame 
placed  some  little  distance  below  it.  Considerable  action  takes  place, 
the  contents  of  the  crucible  undergoing  incipient  deflagration,  and 
becoming  liquid.  It  is  kept  in  a  state  of  fusion  for  a  few  minutes, 
and  allowed  to  cool.  The  crucible  is  then  placed  upon  its  side  in 
a  beaker,  along  with  the  lid,  and  water  cautiously  added,  the  beaker 
being  covered  to  prevent  loss  during  the  effervescence  which  takes 
place  owing  to  the  escape  of  oxygen  from  the  excess  of  peroxide. 

When  the  solid  is  detached  from  the  crucible,  the  latter  is 
lifted  out  with  a  clean  pair  of  tongs  and  rinsed  with  water  ;  the 
lid  is  removed  and  washed  into  the  beaker  in  the  same  way.  The 
solution  is  then  acidified  with  hydrochloric  acid,  filtered  if  neces- 
sary, and  the  sulphuric  acid  precipitated  as  barium  sulphate  by 
means  of  barium  chloride  in  the  usual  manner  (p.  258).$  From 

*  If  there  is  no  manganese  in  the  ore,  the  iron  and  zinc  may  be  separated 
by  precipitation  of  the  former  by  means  of  ammonia,  according  to  the  method 
described  on  p.  277. 

f  Commercial  sodium  peroxide  is  liable  to  contain  traces  of  metallic  sodium, 
in  which  case  a  nickel  crucible  should  be  used.  By  fusing  a  small  quantity  of 
the  peroxide  on  a  scrap  of  platinum  foil  it  is  easy  to  test  whether  or  not  it 
would  be  safe  to  employ  a  platinum  vessel. 

J  Instead  of  employing  sodium  peroxide  as  the  oxidising  agent,  strong 
nitric  acid  may  be  used.  For  this  purpose  the  powdered  ore  is  digested  in  a 
flask  at  a  gentle  heat  (see  Fig.  39)  with  strong  nitric  acid,  to  which,  as  the 
operation  proceeds,  a  little  hydrochloric  acid  or  a  few  crystals  of  potassium 
chlorate  are  added;  or  fuming  nitric  acid  (sp.  gr.  1-5)  may  be  employed 
alone.  When  the  action  is  complete,  and  the  whole  of  the  sulphur  (which 
at  first  often  separates  out  and  floats  about  in  liquid  drops)  is  oxidised  into 
sulphuric  acid,  a  little  pure  sodium  chloride  is  added.  This  converts  the 
sulphuric  acid  into  sodium  sulphate,  and  so  prevents  its  volatilisation  during 
the  subsequent  process  of  evaporation.  The  solution  is  then  evaporated  in  a 
porcelain  dish  upon  a  steam-bath.  When  nearly  dry,  a  little  strong  hydro- 
chloric acid  is  added,  and  the  mixture  evaporated  to  dryness  The  residue 


Analysis  of  Zinc-blende.  281 

the  weight  of  barium  sulphate  obtained,  the  proportion  of  sulphur 
is  calculated.     Factor — 

(BaS04)233  :  (8)32=  1:0-1375 

Silicious  Matter  and  Lead,— About  2  grams  of  the  dry 
and  finely  powdered  ore  are  weighed  out  into  a  beaker,  and  20  to 
25  c.c.  of  strong  hydrochloric  acid  added.  The  beaker  should  be 
immediately  covered,  since  if  any  carbonate  is  present  there  will 
be  effervescence.  The  mixture  is  then  gently  boiled  until  sul- 
phuretted hydrogen  is  no  longer  given  off,  when  about  20  c.c.  of 
strong  nitric  acid  are  added,  and  the  heating  continued  until  the 
decomposition  of  the  ore  is  complete.  The  contents  of  the  beaker 
are  then  rinsed  into  a  porcelain  dish  and  evaporated  to  dryness 
upon  a  steam-bath,  after  which  the  dish  with  the  dry  residue  is 
placed  in  an  air-oven  and  heated  to  about  160°  to  render  the  silica 
insoluble.  The  residue  is  next  moistened  with  strong  hydrochloric 
acid,  and  about  5  c.c.  of  strong  sulphuric  acid,  previously  diluted 
with  twice  its  own  volume  of  water,  added.  The  mixture  is  then 
cautiously  heated  on  a  sand-bath  until  fumes  of  sulphuric  acid  are 
given  off,  by  which  time  all  the  nitrates  and  chlorides  will  have 
been  converted  into  sulphates,  and  their  acids  expelled.  The  dish 
is  then  allowed  to  cool,  water  is  added,  and  the  mixture  filtered. 
The  residue,  consisting  of  silicious  matter  and  lead  sulphate,  is 
washed  with  water  acidulated  with  sulphuric  acid,  the  filtrate  and 
washings  being  set  aside  for  the  estimation  of  the  remaining 
constituents. 

In  order  that  the  residue  may  be  dried  without  the  paper  be- 
coming charred,  the  dilute  acid  is  first  washed  out  of  it  by  means 
of  alcohol,  the  washings  being  neglected.  The  dried  residue  is 
then  transferred  to  a  porcelain  crucible,  and  the  incineration  of  the 
paper  conducted  as  described  in  the  estimation  of  lead  as  sulphate 
(p.  242).  The  weight  obtained  is  that  of  the  silicious  matter  and  the 

is  again  treated  with  hydrochloric  acid,  and  once  more  taken  down  to  dryness 
in  order  to  completely  expel  the  nitric  acid.  It  is  finally  dissolved  in  water 
with  a  little  hydrochloric  acid,  filtered,  and  diluted  to  about  100  c.c.  The 
sulphuric  acid  is  then  precipitated  as  barium  sulphate  in  the  usual  way. 
The  residue  may  contain  lead  sulphate  along  with  the  gangue.  If,  therefore, 
lead  is  present  in  the  ore,  this  residue  must  be  treated  as  described  above, 
and  the  amount  of  sulphur  which  has  been  thus  carried  down  in  combination 
with  the  lead  must  be  added  to  the  total  weight  of  sulphur. 

The  sulphur  in  natural  sulphides  is  usually  estimated  either  by  this  process 
or  by  oxidation  with  sodium  peroxide,  the  latter  being  the  quicker  process. 
Oxidation  of  the  sulphur  may,  however,  be  effected  by  other  methods,  such 
as  by  fusion  with  a  mixture  of  sodium  carbonate  and  potassium  nitrate,  or  by 
digesting  the  ore  in  a  strong  solution  of  caustic  potash,  and  passing  a  stream 
of  chlorine  through  the  mixture. 


282     •          Typical  Gravimetric  Separations. 

lead  sulphate  together.  The  contents  of  the  crucible  are  then 
transferred  to  a  small  beaker,  and  gently  boiled  with  a  solution  of 
ammonium  acetate  and  ammonia  ;  it  is  allowed  to  settle,  and  the 
liquid  decanted  through  a  small  filter.  It  is  treated  to  succes- 
sive extractions  with  fresh  portions  of  the  ammonium  acetate 
solution  until  the  whole  of  the  lead  sulphate  is  dissolved  out,  the 
completion  of  the  process  being  ascertained  by  testing  a  drop  of 
the  filtrate  with  ammonium  sulphide. 

The  residue  is  then  dried,  and  the  paper  incinerated  either  in  the 
crucible  along  with  the  residue,  or  upon  a  platinum  spiral.  The  result 
gives  the  weight  of  the  silicious  matter  alone,  the  difference  between 
this  and  the  former  weighing  being  the  proportion  of  lead  sulphate, 
from  which  the  percentage  of  lead  is  deduced  by  the  factor  on  p.  242. 

Copper  and  Cadmium.— -To  the  solution  containing  the  re- 
maining metals,*  10  to  15  c.c.  of  strong  hydrochloric  acid  are 
added,  and  sulphuretted  hydrogen  is  passed  through  until  precipita- 
tion is  complete.  If  cadmium  is  absent,  the  precipitate  is  treated 
as  described  on  p.  236  (Estimation  of  copper  as  sulphide).  If,  how- 
ever, the  ore  contains  cadmium,  the  precipitate  of  the  mixed  sul- 
phides, after  being  thoroughly  washed,  is  boiled  with  dilute  sulphuric 
acid  (strong  acid  I  part,  water  5  parts).  It  is  then  filtered,  and  the 
insoluble  copper  sulphide  washed  and  dried,  and  treated  as  described 
on  p.  237.  The  filtrate  and  washings  containing  the  cadmium  are 
neutralised  with  ammonia,  one  or  two  drops  of  hydrochloric  acid 
added,  and  the  cadmium  precipitated  as  sulphide  by  means  of  sulphu- 
retted hydrogen.  The  precipitate  is  treated  as  described  on  p.  257. 

Iron.f — The  filtrate  from  the  first  precipitation  with  sulphuretted 
hydrogen  is  boiled  until  the  gas  is  entirely  expelled,  and  a  few 
drops  of  nitric  acid  are  added  in  order  to  re-oxidise  the  iron. 
Barium  chloride  is  then  added  so  long  as  a  precipitate  is  produced, 
and  the  barium  sulphate  filtered  off  and  washed.  The  precipitate 
is  neglected.  In  this  way  the  sulphates  present  are  converted  into 
chlorides.}  The  solution  is  then  nearly  neutralised  with  ammonium 
carbonate,  by  adding  the  reagent  cautiously  until  a  slight  precipitate 
persists,  and  then  adding  hydrochloric  acid  drop  by  drop  with 

*  This  solution  should  not  be  less  than  about  150  c.c.  in  volume.  If  the 
washings  from  the  insoluble  residue  do  not  bring  it  up  to  this,  it  must  be 
diluted  with  water, 

f  If  the  ore  contains  no  manganese,  the  iron  may  be  precipitated  as 
hydroxide  by  means  of  ammonia  (as  described  on  p.  230)  after  the  sulphuretted 
hydrogen  has  been  boiled  off  and  the  iron  re-oxidised  with  nitric  acid. 

j  This  step  is  necessary  to  prevent  the  precipitation  of  basic  ferric  sulphate 
along  with  the  acetate.  In  analyses  where  iron  is  separately  determined  (volume- 
trically  for  example)  and  it  is  only  required  at  this  stage  to  remove  it  before  precipi- 
tating the  manganese,  it  is  not  necessnry  to  convert  the  sulphates  into  chlorides. 


Analysis  of  Zinc-blende,  283 

constant  stirring,  until  the  precipitate  is  almost  completely  re- 
dissolved,  leaving  the  solution  just  slightly  turbid.*  A  strong 
solution  of  ammonium  acetate,  acidified  with  acetic  acid,  is  then 
added,  and  the  solution,  which  assumes  a  red  colour,  is  boiled  for  a 
few  minutes,  when  the  basic  ferric  acetate  separates  as  a  voluminous 
reddish  precipitate.  This  is  allowed  to  settle,  and  the  colourless 
liquid  decanted  through  a  filter,  to  which  the  precipitate  is  then 
transferred  and  washed  with  hot  water.f  The  precipitate  is  then 
dissolved  in  hydrochloric  acid,  the  solution  again  neutralised  with 
ammonium  carbonate,  and  the  iron  reprecipitated  with  ammonium 
acetate,  filtered,  and  washed.  The  filtrate  and  washings  from  this 
second  precipitation  (which  is  necessary  for  the  complete  separation 
of  the  manganese)  are  added  to  those  from  the  first  operation.  The 
precipitate  is  dried,  and  the  filter  incinerated  with  the  precipitate, 
which  is  then  strongly  heated  in  the  crucible  until  it  is  completely 
converted  into  the  sesquioxide,  Fe2O3,  in  which  form  it  is  weighed.} 

Manganese. § — The  filtrates  from  the  double  precipitation  of 
the  iron  are  together  evaporated  down  until  the  volume  is  reduced 
to  about  loo  c.c.  Ammonia  is  then  added  until  the  solution  is 
strongly  alkaline,  and  the  mixture  heated  nearly  to  boiling.  Satu- 
rated bromine  water  is  then  added  to  the  boiling  liquid  until  the 
precipitation  of  the  hydrated  peroxide  is  complete,  and  the  solution 
is  distinctly  yellow. ||  The  precipitate  is  then  filtered  off,  and  washed 
and  dried.  The  dried  precipitate,  after  the  incineration  of  the 
paper,  is  strongly  heated  in  the  crucible  until  it  is  completely  con- 
verted into  the  tetroxide,  Mn3O4,  in  which  form  it  is  weighed.  \ 

Zinc. — The  alkaline  filtrate  from  the  manganese  separation  is 

*  The  least  excess  of  hydrochloric  acid  must  be  avoided. 

t  This  precipitate  is  difficult  to  wash,  especially  if  it  has  been  boiled  too 
long,  when  it  is  apt  to  become  slimy.  In  cases  where  there  is  a  considerable 
quantity  of  the  precipitate,  the  funnel  should  be  kept  hot  by  one  of  the  arrange- 
ments described  on  p.  210. 

J  When  sodium  acetate  is  used  for  the  precipitation  instead  of  the  ammo- 
nium salt,  the  basic  ferric  acetate  must  be  dissolved  in  hydrochloric  acid,  and 
the  iron  precipitated  as  hydroxide  by  means  of  ammonia. 

§  The  manganese  cannot  satisfactorily  be  precipitated  as  the  carbonate  from 
a  solution  containing  considerable  quantities  of  acetates. 

|l  A  useful  alternative  method  consists  in  adding  bromine  itself  to  the  cold 
evaporated  liquid  until  the  solution  is  brown,  and  then  20  to  30  c.c.  of  strong 
ammonia,  after  which  the  liquid  is  warmed. 

If  If  sodium  acetate  has  been  employed  in  the  previous  separation  of  the 
iron  instead  of  ammonium  acetate,  the  process  for  the  estimation  of  the 
manganese  must  be  modified,  since  the  hydrated  peroxide  retains  portions  of 
the  alkali  which  are  not  removed  by  washing,  and  being  non-volatile  they 
remain  behind  with  the  manganese  in  the  crucible.  In  this  case  the  hydrated 
peroxide  must  be  dissolved  in  hydrochloric  acid,  and  the  manganese  precipitated 
as  carbonate  by  means  of  sodium  carbonate  (p.  245). 


284  Typical  Gravimetric  Separations. 

saturated  with  sulphuretted  hydrogen  (or  colourless  ammonium 
sulphide  may  be  added),  and  the  zinc  estimated  as  sulphide,  as 
described  on  p.  244.  Or  the  precipitated  zinc  sulphide,  after  being 
filtered  off,  may  be  dissolved  in  dilute  hydrochloric  acid,  the  solution 
boiled,  and  the  zinc  precipitated  as  carbonate  by  means  of  sodium 
carbonate,  and  estimated  in  the  form  of  zinc  oxide  according  to 
the  method  given  on  p.  243,  or  by  the  volumetric  method  described 
on  p.  371. 

NOTE. — The  analyses  of  other  natural  sulphides,  such  as  copper- 
pyrites,  iron-pyrites,  galena,  etc.,  are  conducted  on  the  same  general 
lines  as  are  described  above,  except  that  certain  necessary  modifi- 
cations have  to  be  introduced  depending  upon  the  particular  metals 
present  in  the  ore.  Thus,  copper  and  iron  pyrites  are  liable  to 
contain  arsenic,  or  arsenic  and  antimony.  These  are  precipitated 
as  sulphides  along  with  the  copper,  and  afterwards  separated  from 
the  copper  sulphide  by  solution  in  sodium  sulphide.*  From  this 
solution  the  sulphides  of  arsenic  and  antimony  are  reprecipitated 
by  acidifying  with  hydrochloric  acid  and  passing  sulphuretted 
hydrogen,  and  the  elements  separated  by  the  following  method. 

Separation  of  Arsenic  and  Antimony. — The  tsvo  sulphides 
are  dissolved  in  strong  hydrochloric  acid,  with  the  addition  of 
crystals  of  potassium  chlorate ;  tartaric  acid  and  ammonium 
chloride  are  added  in  considerable  quantity,  then  an  excess  of 
ammonia.  (The  addition  of  ammonia  should  cause  no  precipitation  ; 
should  a  precipitate  form,  it  is  due  to  insufficient  tartaric  acid  and 
ammonium  chloride,  and  more  must  be  added.)  The  arsenic  is 
then  precipitated  as  ammonium  magnesium  arsenate  f  by  means  of 
"  magnesia  mixture,"  and  weighed  as  magnesium  pyro-arsenate,  as 
described  on  p.  254.  The  antimony  contained  in  the  solution  is 
precipitated  as  antimonious  sulphide  by  means  of  sulphuretted 
hydrogen,  and  finally  weighed  in  the  form  of  the  tetroxide,  as 
described  on  p.  256. 

Galena  often  contains  silver  besides  other  metals.  The  ore  in 
this  case  is  decomposed  by  strong  nitric  acid  (sp.  gr.  1-5),  the  lead 
converted  into  sulphate  by  means  of  sulphuric  acid,  and  the  lead 
sulphate  and  silicious  matter  separated  together.  The  silver  is 
then  precipitated  as  silver  chloride  by  means  of  hydrochloric  acid, 
and  estimated  as  directed  on  p.  238,  while  the  remaining  metals, 
such  as  copper,  antimony,  iron,  zinc,  which  are  liable  to  be  present, 
are  separated  as  explained  above. 

*  Sodium  sulphide  is  prepared  by  saturating  a  solution  of  sodium  hydroxide 
with  sulphuretted  hydrogen,  and  then  adding  fresh  sodium  hydroxide  solution 
until  the  mixture  ceases  to  smell  of  the  gas  (see  "  Reagents,"  Appendix). 

f  Precipitated  under  these  conditions,  especially  if  the  quantity  of  arsenic  is 
large,  the  ammonium  magnesium  arsenate  is  liable  to  contain  a  little  basic 
magnesium  tartrate.  It  is  desirable,  therefore,  to  dissolve  the  precipita;e  in 
hydrochloric  acid,  and  reprecipitate  it  by  the  addition  of  ammonia  and  the 
magnesia  solution. 


Analysis  of  an  Insoluble  Silicate.  285 

Analysis  of  an  Insoluble  Silicate.* 

(Containing  the  metals  iron,  aluminium,  calcium,  magnesium, 
potassium,  and  sodium?) 

Epitome  of  Process. — The  finely  powdered  silicate  is  fused 
with  alkaline  carbonates  (fusion  mixture}.  The  "  melt "  is  extracted 
with  water,  and  the  silicic  acid  precipitated  with  hydrochloric 
acid,  after  which  the  mixture  is  evaporated  to  dryness  and  gently 
heated.  The  insoluble  silica  is  filtered  off,  dried,  and  weighed.  The 
iron  and  aluminium  are  together  precipitated  as  hydroxides,  and 
afterwards  separately  determined  by  one  of  the  methods  given  on 
p.  278. 

The  calcium  is  next  separated  as  calcium  oxalate.  The  filtrate 
is  evaporated  to  dryness,  and  heated  to  expel  ammonium  salts  ;  the 
residue  is  dissolved  in  water,  and  the  magnesium  precipitated  as 
ammonium  magnesium  phosphate. 

The  alkalies  are  estimated  in  a  separate  portion  by  fusion  with 
calcium  carbonate  and  ammonium  chloride,  whereby  they  are  con- 
verted into  soluble  chlorides.  The  product  is  extracted  with  water  ; 
the  soluble  calcium  salts  are  precipitated  as  oxalate,  and  after  the 
removal  of  ammoniacal  salts,  the  mixed  chlorides  of  potassium  and 
sodium  are  dried  and  weighed.  The  proportion  of  potassium 
chloride  in  this  residue  is  then  determined  as  the  double  potassium 
platinum  chloride,  and  the  sodium  chloride  estimated  by  difference  ; 
and  from  the  results  so  obtained  the  percentage  of  potassium  and 
sodium  oxides  is  calculated. 

Silica. — From  1*5  to  3  grams  f  of  the  silicate  (which  has  been 
reduced  to  the  finest  possible  powder,J  and  dried  §),  are  weighed  out 
into  a  platinum  crucible  of  fairly  large  dimensions,  and  intimately 
mixed  with  five  or  six  times  its  weight  of  fusion  mixture  by  means 

*  A  large  number  of  the  common  natural  silicates  (e.g.  the  felspars,  micas, 
hornblende,  etc.)  consist  essentially  of  varying  proportions  of  silicates  of  iron 
and  aluminium,  calcium,  and  magnesium,  and  one  or  more  of  the  alkali  metals. 
The  various  kinds  of  glass  are  also  silicates  containing  similar  constituents, 
being  composed  mainly  of  silicates  of  sodium  (or  potassium)  and  calcium, 
mixed  with  varying  smaller  quantities  of  aluminium,  iron,  and  manganese.  In 
flint  glass  the  calcium  silicate  is  more  or  less  entirely  replaced  by  lead  silicate. 

t  In  the  case  of  such  a  silicate  as  felspar,  which  may  contain  less  than  i  per 
cent,  of  some  bases,  a  larger  quantity  must  be  employed  for  the  analysis  than 
would  be  necessary  in  the  case  of  a  substance  such  as  glass. 

£  It  is  absolutely  necessary,  in  order  to  ensure  the  complete  decomposition 
of  the  silicate,  that  it  should  be  extremely  finely  powdered.  There  should  be 
no  grittiness  to  the  touch  when  it  is  rubbed  between  the  thumb  and  ringers, 
and  the  whole  sample  should  pass  readily  through  a  sieve  of  fine  muslin. 

§  The  powdered  silicate  should  be  freed  from  any  adherent  moisture  by 
being  heated  in  the  steam-oven  before  being  weighed  out  for  analysis.  Some 
silicates  contain  combined  water  (e.g.  serpentine,  meerschaum,  etc. ).  This  should 
be  estimated  by  heating  a  weighed  quantity  of  the  powdered  mineral  to  redness 
in  a  platinum  crucible  until  the  weight  is  constant. 


286  Typical  Gravimetric  Separations. 

of  a  stout  platinum  wire,  or  a  thin  glass  rod  with  carefully  rounded 
ends.  The  entire  mixture  should  not  more  than  half  fill  the  crucible. 
The  covered  crucible  is  then  heated  over  a  Bunsen  flame,  at 
first  gently  in  order  to  expel  the  moisture  present  in  the  mixture, 
and  afterwards  more  strongly  until  the  mass  begins  to  melt  round 
the  edges.  It  is  then  heated  by  means  of  a  blowpipe,  until  the 
decomposition  is  complete  and  the  contents  of  the  crucible  are  in 
a  state  of  quiet  fusion.  Care  must  be  taken,  by  regulating  the  heat 
of  the  blowpipe,  to  prevent  undue  frothing  of  the  mass  while  the 
carbon  dioxide  is  being  evolved  ;  the  progress  of  the  operation 
should  be  watched  by  momentarily  slightly  raising  the  crucible  lid 
from  time  to  time. 

When  the  operation  is  complete,  the  crucible  is  allowed  to  cool 
down,  and  when  cold  it  is  placed  upon  its  side  in  a  beaker  with 
about  100  c.c.  of  water,  care  being  taken  that  no  impurities  are 
conveyed  into  the  solution  upon  the  outside  of  the  crucible.  The 
beaker  is  heated  upon  an  iron  plate  or  sand-bath,  and  the  water 
allowed  to  boil  gently  until  the  "  melt "  is  either  entirely  detached 
from  the  crucible  or  has  become  honeycombed  by  the  action  of  the 
hot  water.  Hydrochloric  acid  is  then  cautiously  added  in  small 
portions  at  a  time  (the  clock-glass  cover  being  partially  withdrawn 
for  each  addition)  until  effervescence  ceases,  and  no  further  pre- 
cipitation of  gelatinous  silicic  acid  takes  place.  The  crucible  and 
lid  are  then  withdrawn  *  and  rinsed  into  the  beaker,  f 

The  mixture  is  transferred  to  a  dish,  and  evaporated  to  dry- 
ness  upon  a  steam-bath,  the  gelatinous  mass,  as  it  stiffens,  being 
stirred  at  frequent  intervals  with  a  short  glass  rod,  in  order  to 
break  it  up  as  much  as  possible  and  thus  expedite  the  drying. 
When  the  mass  has  become  white  and  pulverulent,  the  dish  is  trans- 
ferred to  an  air-bath,  and  heated  to  about  160°  for  half  an  hour. 

The  residue  is  then  moistened  with  a  little  strong  hydrochloric 
acid,  and  digested  upon  the  steam-bath  for  a  short  time,  acid  being 
added  as  evaporation  goes  on.J  Hot  water  is  'added,  and  the  silica 

*  Ordinary  brass  or  iron  crucible  tongs  must  not  be  dipped  into  this  acid 
liquid.  In  the  absence  of  platinum  or  bone-tipped  tongs,  the  crucible  can  be 
lifted  a  little  way  out  of  the  liquid  by  means  of  a  glass  rod,  and  the  projecting 
part  first  rinsed  with  a  drop  of  water  from  the  wash-bottle,  after  which  the 
crucible  can  be  taken  hold  of  with  the  tongs. 

f  If  the  fusion  has  been  satisfactorily  conducted,  there  should  be  no  trace 
of  gritty  particles  at  the  bottom  of  the  beaker.  Should  such  be  detected,  it 
will  probably  be  due  to  imperfect  powdering  of  the  mineral,  and  in  that  case 
the  experiment  must  be  set  aside,  and  the  operation  repeated. 

J  Oxides  of  iron  and  aluminium,  after  being  heated,  are  less  easily  dissolved 
by  hydrochloric  acid,  hence  the  necessity  for  this  step. 


Analysis  of  an  Insoluble  Silicate.  287 

washed  several  times  by  decantation  with  hot  water,  after  which  it 
is  transferred  to  the  filter  and  washed  until  the  wash-water  is  free 
from  chlorides. 

The  silica  is  dried  in  the  steam-oven,  transferred  to  a  platinum 
crucible,*  and  the  paper  incinerated  upon  a  platinum  spiral.  The 
covered  crucible  is  heated,  at  first  very  cautiously  with  a  small 
flame,  and  afterwards  to  a  red  heat,  and  weighed  until  constant.! 

Iron  and  aluminium  are  precipitated  together  in  the  form  of 
hydroxides  by  the  addition  of  ammonium  chloride  and  ammonia  to 
the  filtrate  from  the  silica.  The  precipitation  is  conducted  as  de- 
scribed on  p.  277,  and  the  iron  and  aluminium  are  separately  esti- 
mated by  either  of  the  methods  given  on  p.  278.  From  the  results 
obtained,  the  percentages  of  ferric  oxide,  Fe2O3,  and  aluminium 
sesquioxide,  A12O3,  present  in  the  mineral  are  calculated. 

Calcium* — If  the  filtrate  and  washings  from  the  precipitated 
hydroxides  of  iron  and  aluminium  are  more  in  volume  than 
about  150  c.c.,  they  should  be  evaporated  down  to  that  bulk, 
and  the  calcium  then  precipitated  as  oxalate,  as  on  p.  231.  From 
the  result  obtained,  the  percentage  of  calcium  oxide,  CaO,  is 
calculated. 

Magnesium.—  In  the  filtrate  from  the  calcium  oxalate  the 
magnesium  is  precipitated  in  the  form  of  ammonium  magnesium 
phosphate,  as  described  on  p.  233. 

From  the  weight  of  magnesium  pyrophosphate  obtained,  the 
percentage  of  magnesium  oxide,  MgO,  is  calculated.  Factor — 

(Mg2P207)  222  :  (2MgO)  80  =  i  :  0-3603 

Potassium  and  Sodium — When  the  silicate  has  been 
"opened  up  "  by  fusion  with  alkaline  carbonates  as  above  described, 
it  is  obvious  that  the  estimation  of  the  alkalies  cannot  be  made  in 
the  solution  so  obtained  j  these  constituents  must  therefore  be 

*  The  silica  obtained  in  this  way  is  an  extremely  light  powder,  which  is 
easily  blown  away ;  hence  some  care  is  necessary  in  the  manipulation  of  it,  in 
order  to  avoid  loss  from  this  cause. 

f  It  is  always  desirable  to  test  the  purity  of  the  silica  after  it  has  been 
weighed.  In  the  case  of  silicates  which  are  liable  to  contain  small  quantities 
of  tin,  tungsten,  niobium,  or  tantalum,  the  oxides  of  these  metals  will  be  ad- 
mixed with  the  silica,  and  in  such  cases  an  examination  of  the  silica  is  necessary. 
For  this  purpose  the  contents  of  the  platinum  crucible  are  digested  upon  a 
steam-bath  with  pure  aqueous  hydrofluoric  acid  and  a  few  drops  of  strong 
sulphuric  acid,  the  operation  being  performed  in  a  draught  cupboard.  The 
silica  is  thus  converted  into  silicon  fluoride,  which  is  expelled.  Fresh  hydrofluoric 
acid  is  added  once  or  twice  as  the  liquid  evaporates,  after  which  the  contents 
of  the  crucible  are  evaporated  to  dryness,  strongly  heated,  and  weighed.  The 
residue  is  again  submitted  to  the  same  treatment  until  the  weight  is  constant. 
The  loss  of  weight  represents  the  silica  which  was  present. 


288 


Typical  Gravimetric  Separations. 


determined  in  a  separate  portion  of  the  mineral,  which  is  decomposed 
by  another  method.  About  1*5  to  2  grams  of  the  powdered  silicate 
are  weighed  into  a  platinum  crucible,  and  there  intimately  mixed 
with  about  six  times  its  weight  of  pure  precipitated  calcium 
carbonate,*  and  about  its  own  weight  of  pure  ammonium  chloride. 
The  crucible  is  then  gradually  raised  to  a  bright  red  heat,  and 
maintained  at  that  temperature  for  an  hour.  This  may  be  accom- 
plished by  means  of  a  powerful  Bunsen, 
a  plumbago  crucible  with  the  bottom 
out  being  inverted  over  the  platinum 
crucible,  as  shown  in  Fig.  6o.t  The 
crucible  is  afterwards  allowed  to  cool, 
and  then  placed  in  a  covered  porcelain 
dish  and  digested  with  water  upon  a 
sand-bath.  The  crucible  is  withdrawn 
and  rinsed,  and  the  liquid  filtered,  the 
residue  being  thoroughly  washed.  The 
solution,  which  now  contains  the  alkali 
metals  in  the  form  of  chlorides,  is  freed 
from  any  lime  salts  which  have  dis- 
solved, by  the  addition  of  ammonia, 
ammonium  carbonate,  and  a  small 
quantity  of  ammonium  oxalate.  The 
precipitate  is  filtered  off  and  washed, 
and  the  filtrate  evaporated  to  dryness 
in  a  platinum  dish  upon  a  steam-bath.  The  residue  in  the  dish 
is  heated  over  a  small  rose  burner  without  raising  the  temperature 
to  redness,  in  order  to  expel  all  the  ammonium  salts.  It  is  then 
dissolved  in  a  small  quantity  of  water,  and  the  last  traces  of  lime 
which  may  have  escaped  precipitation  are  thrown  down  by  the 
addition  of  one  or  two  drops  of  ammonia  and  ammonium  oxalate. 
The  solution  is  passed  through  a  small  filter,  which  must  be  after- 
wards thoroughly  washed,  the  filtrate  and  washings  being  received 
in  a  weighed  platinum  dish.  A  few  drops  of  hydrochloric  acid  are 
added,  and  the  liquid  is  then  evaporated  to  dryness,  heated  to 
expel  ammonia,  and  weighed.  It  is  again  heated  and  weighed,  until 
the  weight  of  the  mixed  chlorides  of  the  alkali  metals  is  constant. 

*  Prepared  by  the  addition  of  ammonium  carbonate  to  a  solution  of  barium 
chloride,  and  thoroughly  washing  and  drying  the  precipitate. 

f  The  blowpipe  is  not  so  suitable  for  the  purpose,  partly  on  account  of  the 
tediousness  of  using  a  blowpipe  for  so  long,  but  more  particularly  because  of 
the  liability  to  overheat  portions  of  the  mixture,  and  the  subsequent  risk  of  loss 
of  alkaline  chlorides  by  volatilisation. 


FIG.  60. 


Estimation  of  Alkalies.  289 

In  order  to  estimate  the  relative  proportions  of  the  potassium 
and  sodium  chlorides,  the  residue  is  dissolved  in  a  very  small 
quantity  of  water,  and  the  potassium  precipitated  and  weighed  as 
the  double  potassium  platinic  chloride,  as  described  on  p.  249. 
From  the  weight  of  the  double  salt  obtained,  the  weight  of  potassium 
chloride  is  calculated.  Factor — 

(2KCl,PtCl4)  486  :  (2KC1)  149  =  i  :  0-3065 

And  on  deducting  the  weight  of  potassium  chloride  so  obtained, 
from  the  weight  of  the  mixed  chlorides,  the  proportion  of  sodium 
chloride  is  found. 

Having  found  the  weights  of  the  two  separate  chlorides,  the 
percentages  of  potassium  and  sodium  oxides  (K2O  and  Na2O) 
which  they  represent  is  calculated  by  means  of  the  factors — 

(2KC1)  149  :  (K2O)  94  =  I  :  0-6308 
(2NaCl)  117  :  (Na2O)  62  =  i  :  0-530 

NOTE  i.  Alternative  Methods  for  the  Estimation  of  the 
Alkalies. — (i)  In  the  foregoing  method,  it  will  be  obvious  that,  as 
the  potassium  alone  is  determined,  the  sodium  being  estimated  by 
difference,  the  experimental  errors  fall  more  heavily  upon  the  latter 
alkali.  This  may  be  obviated  by  determining  the  chlorine  (instead 
of  the  potassium)  in  the  mixed  chlorides.  The  mixed  chlorides  are 
dried  and  weighed  as  in  the  process  described  above  ;  the  residue 
is  then  dissolved  in  water,  and  the  chlorine  precipitated  as  silver 
chloride  by  means  of  silver  nitrate.  The  estimation  may  be  carried 
out  gravimetrically,  as  described  on  p.  238,  or  by  the  volumetric 
method  given  on  p.  360. 

From  the  result,  by  whichever  method  obtained,  the  percentage 
of  chlorine  contained  in  the  mixed  chlorides  is  calculated. 

From  these  data,  namely,  (W)  the  weight  of  the  mixed  chlorides, 
and  (B)  the  percentage  of  chlorine  in  the  mixture,  the  actual 
amount  of  each  chloride  present  can  be  calculated  ;  thus — 

The  theoretical  percentage  of  chlorine  in  sodium  chloride  =  6o'68 
„  „  „  „       potassium  chloride  =  47*65 

The  percentage  of  chlorine,  therefore,  in  any  mixture  of  these 
chlorides  will  obviously  lie  between  these  two  extremes  ;  the  nearer 
it  approaches  to  one  of  them,  the  less  of  the  other  chloride  will  be 
present  in  the  mixture.  Thus,  for  example,  the  percentage  of 
sodium  chloride  in  the  mixture  will  rise  from  o  to  100  as  the  per- 
centage of  chlorine  increases  from  47-65  to  60-68,  i.e.  passes  through 
a  range  of  13*03  per  cents.  (60*68  —  47^65  =  13*03). 

Hence,  if  the  percentage  of  chlorine  in  potassium  chloride  be 
deducted  from  the  percentage  of  chlorine  found  in  the  mixed 
chlorides,  the  result  when  multiplied  by  100  and  divided  by  13-03 
gives  the  percentage  of  sodium  chloride  in  the  mixture— 


290  Typical  Gravimetric  Separations. 

13-03  :  B  —  47-65  : :  100  :  x  =  percentage  of  sodium  chloride  in 

the  mixed  chlorides 

From  this  the  actual  weight  of  sodium  chloride  found  is  calcu- 
lated by  the  proportion — 

loo  :  W  : :  x  :  y  =  grams  of  sodium  chloride  in  the  mixed  chlorides 

And  then — 

W  —  y  —  2  =  grams  of  potassium  chloride  in  the  mixed  chlorides 
A  concrete  example  will  render  this  perfectly  clear — 

The  weight  of  the  mixed  chlorides  (W)  was  0-64  gram 
The  percentage  of  chlorine  in  the  mixture  (B)  was  50*256 

Deducting    from    B   the    theoretical    percentage   of   chlorine    in 
potassium  chloride,  we  get — 

Percentage  of  chlorine  found     50*256 

,,  ,,         in  KC1 47*650 

Difference         2-606 

Then  13-03  :  2'6o6; !  100:20-0  =  °/0  of  NaCl  in  mixed  chlorides 
and  loo  :  0*64  : :  20  :  0-1280  =  grams  of  NaCl  in  mixed  chlorides 
And  deducting  this  from  the  weight  of  the  mixed  chlorides,  the 
weight  of  the  potassium  chloride  is  found  ;  thus — 

Grams. 
Weight  of  mixed  chlorides        ...         ...         ...         ...     0*640 

,,        sodium  chloride  .     0-128 


,,        potassium   ,,  0*512 

By  multiplying  the  weights  of  sodium  chloride  and  of  potassium 
chloride  by  their  respective  factors  (given  above),  the  weights  of  the 
oxides  equivalent  to  the  chlorides  will  be  obtained,  and  from 
these  the  actual  percentages  of  Na2O  and  K2O  in  the  silicate  are 
calculated  ;  thus — 

0*128  X  0*530  =  0*06784  =  grams  Na2O 
0-512  x  0*6308  =  0*32297  =  grams  K2O 
Weight  of  silicate  taken  for  analysis  =  2*0150  grams 
therefore  2-0150  :  0*06784  : :  100  :  3-3  r  =  °/0  of  Na2O  in  the  silicate 
and  2-0150  :  0-32297  ::  loo  :  15*03  =  °/0  K2O          „         „ 

(2)  Another  method  for  the  estimation  of  the  alkali  metals  is  an 
electrolytic  process,*  by  means  of  which  both  alkali  metals  can  be 
separately  determined.  The  mixed  chlorides  are  dried  and  weighed 
as  before,!  and  the  potassium  precipitated  as  the  double  potassium 
platinic  chloride.  The  precipitate  is  washed  perfectly  free  from 

*  See  Section  V.  p.  292. 

f  It  is  not  essential  to  weigh  the  dry  chlorides,  although  a  knowledge 
of  their  weight  serves  as  a  useful  check  upon  the  results  obtained  by  the 
electrolysis. 


Estimation  of  Alkalies.  291 

platinum  chloride  by  means  of  alcohol  ;  it  is  then  dissolved  in 
water  and  the  solution  submitted  to  electrolysis,  using  the  current 
from  one  Daniell  cell.  The  deposited  platinum  is  washed  and 
dried,  and  from  its  weight  the  weight  of  potassium  chloride  present 
is  calculated  by  means  of  the  factor — 

(Pt)  195  :  (2KC1)  149  =  i  :  07641 

The  filtrate  from  the  precipitated  potassium  platinic  chloride 
contains  the  sodium  chloride  mixed  with  the  excess  of  platinum 
chloride.  It  is  electrolysed  until  the  whole  of  the  platinum  is 
deposited,  and  the  liquid,  now  containing  only  sodium  chloride,  is 
transferred  to  a  dish  and  evaporated  to  dryness  ;  after  which  it  is 
gently  heated,  and  weighed.  The  weight  of  sodium  chloride  thus 
determined,  when  added  to  the  weight  of  potassium  chloride  cal- 
culated from  the  former  operation,  should  equal  the  weight  of  the 
mixed  chlorides. 

From  the  chlorides,  the  proportion  of  potash  and  soda  present 
in  the  silicate  is  calculated  as  in  the  previous  example. 

NOTE  2.  Alternative  Method  for  decomposing  the  Silicate ,  by 
means  of  Hydrofluoric  Acid. — The  finely  powdered  silicate  is  treated 
with  strong  aqueous  hydrofluoric  acid  (pure)  in  a  capacious 
platinum  crucible,  and  the  mixture  digested  upon  a  steam-bath 
(in  a  draught  chamber)  for  some  time,  fresh  acid  being  added  once 
or  twice  as  the  liquid  evaporates.  A  few  drops  of  strong  sulphuric 
acid,  previously  mixed  with  its  own  volume  of  water,  are  added  in 
order  to  convert  the  metals  into  sulphates,  and  the  mixture 
evaporated  nearly  to  dryness  upon  the  steam-bath.  The  crucible 
is  then  cautiously  heated  over  a  small  rose  burner  (see  Fig.  52,  p.  248) 
in  order  to  expel  the  excess  of  sulphuric  acid.  The  residue  is 
moistened  with  a  few  drops  of  strong  hydrochloric  acid  and 
dissolved  in  water.  If  the  decomposition  of  the  silicate  by  the 
hydrofluoric  acid  has  been  complete,  any  undissolved  residue  will 
consist  of  insoluble  sulphates  (i.e.  sulphates  of  lead  or  the  alkaline 
earths). 


SECTION   V. 
ELECTROLYTIC  METHODS. 

THE  quantitative  estimation  of  metals  by  the  process  of  electrolysis 
is  based  upon  the  fact  that  when  solutions  of  certain  metallic  salts 
are  submitted  to  the  action  of  an  electric  current  of  suitable 
strength,  the  metals  are  precipitated  upon  the  negative  electrode 
in  the  form  of  coherent  films  or  deposits.  The  operation  is,  in 
fact,  an  electroplating  process,  where  the  article  to  be  "plated" 
is  a  weighed  piece  of  platinum,  and  the  metal  with  which  it  is 
coated  is  the  metal  that  is  to  be  estimated. 

Just  as  in  the  familiar  process  of  electroplating  with  silver,  it 
has  been  found  that  a  solution  of  the  double  cyanide  of  potassium 
and  silver  is  the  most  suitable  compound  for  the  purpose,  so  with 
the  other  metals  it  is  only  certain  of  their  salts  which  are  adapted 
for  the  quantitative  deposition  of  the  metal  they  contain.  For 
example,  when  a  solution  of  zinc  sulphate  is  electrolysed,  the  first 
action  of  the  current  (as  in  other  cases)  is  to  separate  the  salt  into 
its  ions.  The  positive  ions,  or  cathions  (i.e.  the  zinc  atoms),  travel 
to  the  cathode,  or  negative  electrode,  while  the  negative  ions,  or 
anions  (namely,  the  (SO4)  groups),  go  to  the  anode,  or  positive 
electrode;  thus — 

Cathode.       Anode. 

ZnSO4  =  Zn    +    SO4 

The  negative  ions,  however,  at  once  undergo  a  secondary 
decomposition,  for  in  the  presence  of  the  water  they  are  converted 
into  sulphuric  acid  and  free  oxygen — 

2SO4  +  2H2O  =  2H2SO4  +  O2 

The  sulphuric  acid  thus  generated  by  this  secondary  action 
attacks  the  deposited  zinc,  and  a  condition  of  equilibrium  is 
established  when  the  metal  is  dissolved  as  fast  as  it  is  deposited  ; 
hence  complete  precipitation  of  zinc  under  these  conditions  is 
impossible.  If,  instead  of  the  sulphate,  the  oxalate  (or  the  double 
oxalate  of  zinc  and  ammonium,  as  being  the  more  soluble  com- 
pound) be  employed,  no  such  complication  from  the  action  of  the 
products  of  secondary  decompositions  will  arise.  The  zinc  oxalate 


Electrolytic  Methods. 


293 


is  decomposed  into  zinc  and  carbon  dioxide,  while  the  ammonium 
oxalate  is  separated  primarily  into  ammonium  (NH4)  and  carbon 
dioxide,  and  secondarily  into  hydrogen  and  hydrogen  ammonium 
carbonate ;  thus — 


Cathode. 


Anode. 

2C02 
2C02 


ZnC2O4  =  Zn 

(i)  (NH4)2C204  =  2NH4(=  H2  +  2NH3) 
(2)  2NH3  +  2H20  +  2C02  =  2H(NH4)C03 

The  apparatus  required  consists  essentially  of  two  platinum 
electrodes.     That  which  is  to  serve  as  the  cathode,  and  receive 
the  deposited  metal,  is  usually  made  in  the  form  either  of  a  cylinder 
or,  better,  a  cone  ;  which  in  either  case  is 
riveted  or  welded  to  a  stout  platinum  wire, 
a,  Fig.  61.    The  advantage  of  the  conical 
shape  is  that  loss  from  spitting,  as  the 
gas-bubbles  rise  to   the  surface  of  the 
liquid,  is  reduced  to  a  minimum.    The 
anode  is  conveniently  a  thick  platinum 
wire  bent  in  the  form  shown  at  b,  Fig.  61. 

When  only  an  occasional  analysis  is 
to  be  made,  these  electrodes  may  be  sup- 
ported in  a  beaker  by  means  of  ordinary 
clamps  upon  two  retort-stands ;  the  beaker 
being  placed  upon  a  wooden  block,  so  that  by  the  withdrawal  of 
the  latter  the  beaker  can  be  lowered  and  removed.  By  means  of 


FIG.  61 


FIG.  62 


294 


Electrolytic  Methods. 


binding-screws,  the  electrodes  are  connected  to  the  wires  leading 
from  the  battery.     The  arrangement  is  shown  in  Fig.  62. 

Instead  of  conducting  the  electrolysis  in  a  beaker,  it  may  be 
carried  out  in  a  platinum  dish.  In  this  case  the  dish  is  made  the 
cathode,  by  being  supported  either  upon  the  ring  of  a  retort-stand, 
or  upon  any  other  convenient  metal  support  which  is  connected  to 


FIG. 


the  negative  wire  from  the  battery  ;  and  the  anode,  held  in  a  clamp, 
is  lowered  nearly  to  the  bottom  of  the  dish,  as  shown  in  Fig.  63. 
The  metal  to  be  determined  is  then  deposited  upon  the  dish  itself. 
The  deposited  metal  is  ultimately  removed  from  the  electrode  by 
solution  in  a  suitable  acid. 

Typical  Examples.* 

Estimation  of  Copper.—  About  1-5  gram  of  recrystallised 
copper  sulphate  is  weighed  out  into  a  beaker  and  dissolved  in 
200  c.c.  of  water.  About  2  c.c.  strong  nitric  acid  are  added,f  and  the 
mixture  gently  stirred  with  the  platinum  wire  which  is  to  be  used 
as  the  anode  ;  this  is  then  left  standing  in  the  beaker.  The  platinum 

*  Students  who  wish  for  a  more  extended  practice  in  electrolytic  methods, 
should  consult  Classen's  "Quantitative  Analysis  by  Electrolysis,"  English 
translation  (Herrick).  1894. 

t  When  copper  is  to  be  separated  from  other  metals,  a  larger  quantity  of 
nitric  acid  must  be  added  (see  p.  298). 


Estimation  of  Copper.  295 

cone,  which  has  been  perfectly  cleaned  *  and  carefully  weighed,  is 
then  lowered  into  the  beaker  over  the  projecting  end  of  the  anode. 
The  two  electrodes  are  then  supported  by  the  clamps,  so  that  they 
do  not  touch  each  other,  and  are  slightly  raised  from  the  bottom 
of  the  beaker,  the  anode  reaching  slightly  below  the  cone,  in  the 
manner  shown  in  Fig.  62.  • 

The  electrodes  are  then  connected  to  the  battery  (the  platinum 
cone  being  attached  to  the  negative  wire),  and  a  current  of  0*5  to  i 
ampere  passed  through  the  solution.!  The  copper  is  gradually 
deposited  upon  the  platinum  cone,  and  the  blue  colour  of  the  solu- 
tion becomes  fainter  and  fainter,  until  at  last  the  liquid  appears 
colourless.  The  operation  takes  several  hours  for  its  completion, 
and  may  be  conveniently  allowed  to  go  on  all  night.  To  ascertain 
whether  the  precipitation  is  finished,  one  or  two  drops  of  the  solu- 
tion are  withdrawn  by  means  of  a  pipette,  and  a  little  sulphuretted 
hydrogen  added,  which  should  produce  no  coloration.  The 
wooden  block  is  then  withdrawn,  and  the  beaker  lowered  away 
before  interrupting  the  current,  the  anode  at  the  same  time  being 
disconnected  and  removed  with  the  beaker.  The  cone  is  then  dis- 
mounted, and  rinsed  with  water  by  means  of  a  wash-bottle.  It  is 
then  dipped  once  or  twice  into  a  beaker  of  alcohol,  and  placed  in 
the  steam-oven  for  a  few  minutes  to  dry,  and  then  weighed. 

The  gain  in  weight  gives  the  copper  in  the  amount  of  copper 
sulphate  employed,  from  which  the  percentage  of  copper  is  calcu- 
lated. (The  result  is  usually  from  o'i  to  o'2  per  cent,  too  low.) 

The  electrolytic  method  is  one  that  is  much  employed  in  the 
commercial  analysis  or  valuation  of  copper  ores  and  of  metallic 
copper,  as  well  as  in  the  estimation  of  copper  in  alloys.  The 
analyses  are  carried  out  in  the  following  manner  : — 

I.  In  copper  ores. 

*  It  is  of  the  greatest  importance  that  the  cathodes  used  in  electrolytic 
analysis  should  be  absolutely  clean  and  free  from  the  slightest  trace  of  grease, 
even  such  as  would  be  contracted  by  touching  them  with  the  fingers. 

f  When  the  electric  current  from  a  "dynamo"  or  from  storage  batteries 
is  employed,  it  must  be  reduced  to  the  requisite  strength  by  the  introduction  of 
a  system  of  resistance  coils.  If  too  strong  a  current  be  used,  the  deposited 
copper  is  rough  and  less  coherent.  When  only  an  occasional  analysis  is  to  be 
made,  two  or  three  Daniell  cells  may  be  used.  In  this  case  the  zinc  plate  of 
the  battery  is  the  negative,  and  is  connected  to  the  platinum  cone.  Care  must 
be  taken  that  all  the  connections  are  clean,  so  as  to  ensure  perfect  metallic 
contact.  If  the  electrolysis  be  conducted  in  a  platinum  basin,  the  latter  should 
be  covered,  in  order  to  prevent  any  of  the  liquid  being  carried  off  in  the  form 
of  fine  spray  with  the  gas  which  escapes  from  the  anode.  For  this  purpose, 
either  a  clock-glass,  with  a  small  hole  bored  in  the  middle  to  allow  the  wire  to 
project  through,  or  an  inverted  funnel  of  such  a  size  that  the  mouth  will  just  go 
inside  the  dish,  may  be  employed.  When  the  platinum  cone  is  used,  the 
vessel  need  not  be  covered. 


296  Electrolytic  Methods. 

From  i  to  i£  gram  *  of  the  finely  powdered  ore  is  weighed  out 
into  a  porcelain  dish,  and  treated  with  from  10  to  20  c.c.  of  strong 
nitric  acid.  About  an  equal  volume  of  dilute  sulphuric  acid  f  is 
then  added,  and  the  mixture  gently  evaporated  to  about  half  its 
bulk  in  the  covered  dish.  Water  is  then  added,  and  the  insoluble 
residue  (the  gangue)  is  filtered  and  washed.:}:  The  filtrate  is  diluted 
up  to  200  c.c.  with  water,  10  c.c.  of  nitric  acid  added,  and  the  solu- 
tion electrolysed  as  described  above,  with  a  current  from  two  Daniell 
cells.! 

II.  In  commercial  copper. 

About  0-5  gram  of  the  metal  is  weighed  out  into  a  beaker,  and 
dissolved  in  10  c.c.  strong  nitric  acid,  to  which  an  equal  volume  of 
water  has  been  previously  added.  The  mixture  is  boiled  in  the 
covered  beaker  until  "  nitrous  fumes "  cease  to  come  off,  after 
which  it  is  diluted  up  to  200  c.c.  with  water,  and  the  solution 
electrolysed.  The  electrolysis  must  not  be  continued  longer  than 
is  necessary  for  the  complete  precipitation  of  the  copper,  other- 
wise traces  of  antimony  and  arsenic,  which  may  be  present,  will 
also  be  deposited,  which  will  be  evident  by  the  cathode  becoming 
darkened  in  colour.  (For  the  separation  of  traces  of  antimony 
and  arsenic  thus  deposited,  see  footnote.)  Should  the  sample  of 
copper  contain  any  silver,  this  will  be  deposited  and  weighed  along 
with  the  precipitated  copper.  In  this  case  the  silver  must  be 
separately  estimated  (by  dissolving  10  to  20  grams  of  the  metal 
in  nitric  acid,  and  precipitating  with  hydrochloric  acid  in  the  usual 
way,  p.  238)  and  the  proportion  present  deducted  from  the  weight 
obtained. 

Estimation  of  Zinc.— About  2  grams  of  crystallised  zinc 
sulphate  are  weighed  out  into  a  beaker,  and  dissolved  in  about 

*  For  ores  containing  less  than  25  per  cent,  of  copper,  from  ii  to  5  grams 
may  be  taken. 

f  If  the  ore  is  a  sulphide,  the  addition  of  sulphuric  acid  is  unnecessary, 
since  the  sulphur  is  oxidised  into  sulphuric  acid. 

\  Copper  ores  frequently  contain  organic  matter,  in  which  case  this  residue 
will  be  dark  coloured,  and  is  liable  to  retain  a  small  portion  of  the  copper.  If 
the  quantity  of  this  bituminous  matter  is  appreciable,  it  should  be  destroyed 
before  treatment  with  nitric  acid,  by  roasting  the  weighed  quantity  of  powdered 
ore  taken  for  the  analysis  in  a  porcelain  crucible,  with  free  access  of  air. 

§  If  more  than  small  quantities  of  antimony  or  arsenic  are  present  in  the 
ore,  these  elements  begin  to  deposit  upon  the  copper  towards  the  end  of  the 
operation.  When  the  amount  is  quite  small,  they  may  be  separated  by 
strongly  heating  the  electrode.  The  arsenic  volatilises,  while  the  antimony 
and  copper  are  oxidised  ;  the  former  oxide  volatilises,  leaving  copper  oxide. 
This  is  then  dissolved  in  nitric  acid  and  redeposited  by  the  current.  When  the 
amount  of  antimony  or  arsenic  is  considerable,  this  plan  cannot  be  adopted, 
and,  moreover,  the  precipitation  of  copper  in  the  presence  of  much  arsenic  is 
incomplete.  In  this  case,  a  weighed  quantity  of  the  finely  powdered  ore  is 
gently  heated  in  a  covered  porcelain  crucible  with  about  four  times  its  weight 
of  ammonium  chloride.  Antimony  and  arsenic  are  thus  converted  into 
chlorides  and  expelled.  The  residue  is  dissolved  in  nitric  acid  and  treated  as 
already  described  (Classen.  "  Zts  Anal.  Ch.,"  18,  388). 


Analysis  of  German  Silver.  297 

50  c.c.  of  water.  Six  or  7  grams  of  ammonium  oxalate,  dissolved  in 
a  small  quantity  of  warm  water,  are  gradually  added,  with  constant 
stirring. 

The  solution  is  then  diluted  to  150  c.c.  and  electrolysed.  The 
process  is  complete  when  a  drop  of  the  solution  gives  no  precipitate 
when  warmed  with  potassium  ferrocyanide  upon  a  watch-glass. 
The  platinum  electrode  containing  the  deposited  zinc  is  then 
removed,  thoroughly  rinsed  with  water,  and  finally  with  absolute 
alcohol.  It  is  then  placed  for  a  few  minutes  in  a  steam-oven  to 
dry,  and  weighed. 

Estimation  of  Nickel. — About  2  grams  of  ammonium  nickel 
sulphate,  (NH4)2SO4,NiSO4,6H2O,  are  weighed  out  into  a  beaker 
and  dissolved  in  water,  and  the  solution  rendered  strongly  alkaline 
by  the  addition  of  ammonia.  The  volume  of  the  liquid  is  made  up 
to  150  c.c.  with  water,  and  the  solution  submitted  to  electrolysis. 
The  process  may  be  allowed  to  continue  all  night  :  it  is  complete 
when  a  drop  of  the  liquid  gives  no  precipitate  with  ammonium 
sulphide.  The  cathode  is  then  washed  and  dried,  and  finally 
weighed,  as  in  the  former  examples.* 

Analysis  of  Silver  Coin. — About  0-5  gram  of  the  metal  is 
weighed  out  and  dissolved  in  nitric  acid  (strong  acid  diluted  with 
an  equal  volume  of  water)  in  a  covered  evaporating-dish.  The 
liquid  is  evaporated  nearly  to  dry  ness  upon  a  steam-bath  to  expel 
the  acid,  and  the  residue  dissolved  in  water.  The  solution  is  then 
transferred  to  a  beaker,  and  a  moderately  strong  solution  of  ammo- 
nium oxalate  is  added.  Oxalates  of  silver  and  copper  are  formed  ; 
the  latter,  however,  dissolves  in  excess  of  the  reagent,  leaving  the 
insoluble  silver  oxalate  as  a  white  precipitate.  The  liquid  is 
filtered,  and  the  precipitate  washed  perfectly  free  from  copper,  first 
with  a  dilute  solution  of  ammonium  oxalate,  and  finally  with  water- 
The  copper,  contained  in  the  filtrate  and  washings  in  the  form  of 
the  double  ammonium  copper  oxalate,  is  then  precipitated  by  elec- 
trolysis. The  precipitated  silver  oxalate  is  dissolved  in  a  solution 
of  potassium  cyanide,  diluted  with  water  to  about  150  to  200  c.c., 
and  the  solution  electrolysed.  The  electrode  with  the  deposited 
silver  is  washed  with  water  and  with  alcohol,  and  dried  in  the 
steam-oven. 

Analysis  of  German  Silver. — About  i  gram  of  the  alloy  is 

*  Nickel  may  also  be  precipitated  from  the  solution  of  the  double  ammo- 
nium oxalate,  as  described  for  zinc.  In  this  case,  the  solution  should  be 
maintained  at  a  temperature  about  40°  to  50°  throughout  the  operation,  by 
means  of  a  small  flame  placed  beneath  the  beaker.  Cobalt  and  iron  may  be 
estimated  in  the  same  way. 


298  Electrolytic  Methods. 

weighed  out  into  a  beaker,  and  dissolved  in  about  1  5  c.c.  of  strong 
nitric  acid  mixed  with  the  same  volume  of  water.  The  solution  is 
then  diluted  up  to  150  c.c.  with  water,  and  submitted  to  electro- 
lysis. Under  these  circumstances  the  copper  alone  is  deposited.* 
When  the  precipitation  is  complete,!  the  beaker  containing  the 
solution  is  lowered  away  from  the  electrodes  before  the  current 
is  interrupted.  The  cathode  with  its  deposit  of  copper  is  rinsed 
with  water  (the  washing  being  added  to  the  solution  in  the  beaker), 
and  finally  dipped  into  alcohol  and  dried  and  weighed. 

The  positive  electrode  is  also  rinsed  into  the  beaker,$  and  the 
solution  evaporated  to  dryness  in  a  dish  upon  a  water-bath,  with 
the  addition  of  a  little  hydrochloric  acid,  in  order  to  convert  the 
remaining  metals  into  chlorides. 

The  residue  is  dissolved  in  water  and  a  few  drops  of  hydro- 
chloric acid,  and  the  solution  transferred  to  a  beaker.  Sodium 
carbonate  solution  is  added  until  a  slight  precipitate  persists,  after 
which  hydrochloric  acid  is  added  drop  by  drop  until  the  preci- 
pitate just  redissolves.  The  zinc  is  then  precipitated  from  this 
solution  as  zinc  sulphide,  as  described  on  p.  244.  The  washed 
precipitate  is  dissolved  in  the  smallest  possible  quantity  of  strong 
hydrochloric  acid,  and  the  solution  evaporated  to  expel  the  excess 
of  acid.  The  residue  is  dissolved  in  water,  ammonium  oxalate  is 
added  (as  described  on  p.  297),  and  the  solution  diluted  with  water 
to  loo  c.c.  and  electrolysed.  §  , 

*  It  must  be  remembered  that  nitric  acid  is  'decomposed  by  the  passage 
of  an  electric  current,  nitrogen  peroxide  and  oxygen  being  evolved  at  the 
anode,  while  hydrogen  is  liberated  at  the  negative  electrode.  This  nascent 
hydrogen  reacts  upon  the  nitric  acid,  with  the  formation  of  ammonia  ;  thus  — 


HNO3  +  4H2  =  3H20  +  NH3 

Hence,  if  the  current  is  allowed  to  continue  passing  through  the  solution  after 
the  copper  is  all  precipitated,  the  nitric  acid  is  gradually  decomposed  in  this 
manner,  and  then  the  zinc  in  the  solution  begins  to  deposit  along  with  the 
copper. 

t  A  drop  of  the  solution,  when  tested  by  adding  sodium  bicarbonate  and 
potassium  ferrocyanide,  should  give  no  brown  coloration  of  copper  ferro- 
cyanide. 

J  In  the  analysis  of  alloys  in  which  traces  of  lead  are  present  as  an 
impurity,  the  lead  will  be  deposited  as  the  peroxide  upon  the  positive  wire, 
which  will  appear  brownish  or  black  in  consequence.  By  weighing  the 
electrode  with  the  deposited  peroxide,  after  drying  in  an  air-oven  at  no0, 
the  amount  of  lead  present  can  be  calculated. 

§  Any  traces  of  iron  present  in  the  alloy  will  be  deposited  along  with  the 
zinc.  If  the  amount  is  greater  than  mere  traces,  it  may  either  be  removed 
before  the  zinc  is  precipitated  as  sulphide  (as  described  on  p.  271),  or  a 
separate  estimation  may  be  made  in  a  larger  quantity  of  the  alloy  ;  and  the 
calculated  proportion  which  will  have  been  deposited  with  the  zinc  is  deducted 
from  the  weight  obtained  of  the  electrolytic  deposit  of  that  metal. 


Analysis  of  German  Silver.  299 

The  filtrate  and  washings  from  the  zinc  sulphide  contain  the 
nickel,  which  may  be  precipitated  either  from  the  double  ammo- 
nium sulphate  or  oxalate. 

(a)  From  the  Double  Sulphate. — Three  or  4  grams  of  ammo- 
nium sulphate  dissolved  in  10  c.c.  of  water  are  added  to  the 
solution,  and  then  about  20  c.c.  of  strong  ammonia.  The  mixture 
is  diluted  with  water  to  100  c.c.,  and  electrolysed  as  on  p.  297. 

(j8)  From  the  Double  Oxalate. — Four  or  5  grams  of  ammonium 
oxalate,  dissolved  in  about  20  c.c.  of  warm  water,  are  added  to  the 
solution,  which  is  then  diluted  with  water  to  100  c.c.  and  submitted 
to  electrolysis,  the  liquid  being  maintained  at  a  temperature  about 
40°  to  50°  throughout  the  operation. 

Alloys  consisting  of  copper  and  nickel  only,  may  be  analysed 
electrolytically  in  the  following  way  :  From  0*5  to  075  gram  of  the 
alloy  is  dissolved  in  the  minimum  quantity  of  nitric  acid  (diluted 
as  above),  and  the  solution  evaporated  to  dryness  with  the  addition 
of  a  few  drops  of  sulphuric  acid.  The  residue  is  taken  up  with 
water,  and  again  evaporated  to  dryness  with  a  few  drops  of  sul- 
phuric acid  to  completely  expel  the  nitric  acid.  The  residue  is  dis- 
solved in  water  with  a  few  drops  of  sulphuric  acid,  and  made  up  to 
about  icoc.c.,  and  the  solution  electrolysed  (a  current  of  0*5  ampere 
requiring  4  to  6  hours).  The  electrode  is  washed  and  dried  and 
weighed  as  already  described.  To  the  solution  15  c.c.  of  strong 
ammonia  are  added,  and  the  liquid  electrolysed  with  a  current  of 
about  0-3  ampere  (or  three  Daniell  cells).  The  nickel  will  be 
entirely  deposited  in  about  6  hours. 


PART   II. 

VOLUMETRIC  METHODS. 

SECTION  I. 
PRELIMINARY  MANIPULATIONS. 

Introductory  Remarks. — In  volumetric  methods  of  analysis,* 
determinations  are  made,  not  by  weighing  a  product  obtained  by 
the  interaction  of  a  reagent  with  the  substance  to  be  estimated, 
but  by  finding  the  weight  of  the  reagent  which  it  is  necessary  to 
employ  in  order  to  exactly  complete  a  given  chemical  interaction 
with  the  substance  to  be  determined.  For  example,  in  making  a 
gravimetric  determination  of  chlorine  in  a  soluble  chloride,  the 
chlorine  is  precipitated  as  silver  chloride  by  the  addition  of  silver 
nitrate  solution  in  quantity  a  little  in  excess  of  that  which  is  required 
to  complete  the  reaction,  and  the  precipitated  product  of  the  inter- 
action is  washed  and  weighed.  The  volumetric  method,  on  the 
other  hand,  consists  in  finding  the  weight  of  the  silver  nitrate  which 
is  required  in  order  to  just  exactly  throw  down  the  whole  of  the 
chlorine  as  silver  chloride.  Or  again,  instead  of  estimating  iron  by 
precipitating  the  hydroxide,  and  weighing  the  sesquioxide  obtained 
by  heating  this  hydroxide,  the  volumetric  method  consists  in  find- 
ing the  weight  of  some  oxidising  agent,  such  as  potassium  per- 
manganate or  potassium  dichromate,  which  is  just  exactly  required 
to  oxidise  the  iron  from  the  ferrous  to  the  ferric  state. 

The  weight  of  the  reagent  which  is  used  in  a  volumetric  deter- 
mination is  ascertained  by  employing  a  solution  of  known  and 
definite  strength,  and  measuring  the  volume  of  it  which  is  required 
to  carry  the  chemical  change  to  its  completion.  Thus,  in  the  above 
illustration,  if  the  exact  strength  of  the  silver  nitrate  solution  be 

*  The  term  "  volumetric  analysis  "  is  here  employed  in  its  more  restricted 
sense,  as  applying  to  those  methods  of  analysis  where  the  volumes  of  solutions 
only  are  concerned.  Strictly  speaking,  it  also  includes  those  analytical  pro- 
cesses involving  the  measurement  of  volumes  of  gases  or  vapours,  but  by  general 
custom  these  are  classed  together  under  the  head  of  "  gas  analysis,"  and  con- 
veniently constitute  a  separate  section  of  analytical  practice. 


Classification  of  Volumetric  Methods.  301 

known,  it  is  only  necessary  to  carefully  measure  the  volume  which 
must  be  employed  for  the  exact  precipitation  of  the  whole  of  the 
chlorine,  in  order  to  learn  the  weight  of  the  silver  nitrate  used,  and 
from  this  the  weight  of  chlorine  precipitated  can  be  calculated. 

Similarly,  if  the  strength  of  the  potassium  permanganate  solu- 
tion be  known,  then  from  the  number  of  cubic  centimetres  employed, 
the  weight  of  the  permanganate  which  is  required  to  oxidise  the 
iron  present  can  at  once  be  ascertained,  and  from  this  the  actual 
quantity  of  the  iron  is  readily  calculated. 

In  practice  these  calculations  are  made  once  for  all  when  the 
reagents  of  known  definite  strengths,  called  standard  solutions,  are 
prepared  ;  so  that  the  -volume  of  such  a  reagent  which  is  required 
for  the  analysis,  gives  at  once  the  weight  of  the  substance  being 
estimated.  Thus,  in  the  above  examples,  the  strength  of  the 
standard  silver  nitrate  being  known,  the  weight  of  chlorine  which 
can  be  precipitated  by  I  c.c.  is  calculated,  and  this  weight  multiplied 
by  the  number  of  cubic  centimetres  employed  for  an  analysis  gives 
the  weight  of  chlorine  in  the  compound.  Suppose,  for  example, 
that  the  standard  silver  nitrate  contains  ofoi7  gram  of  the  salt  in 
every  cubic  centimetre,  this  would  be  capable  of  precipitating 
0-00355  gram  of  chlorine  ;  hence  I  c.c.  of  the  standard  silver 
nitrate  represents,  or  is  equivalent  to,  0-00355  gram  of  chlorine. 

In  the  same  manner,  from  the  known  strength  of  the  potassium 
permanganate  solution,  the  weight  of  iron  which  i  c.c.  of  it  is 
capable  of  oxidising  can  be  calculated ;  so  that  when  the  reagent 
is  employed  for  the  estimation  of  iron,  the  number  of  cubic  centi- 
metres used,  multiplied  by  the  weight  of  iron  which  each  cubic 
centimetre  is  equivalent  to,  will  give  the  weight  of  iron  in  the  portion 
of  the  substance  being  analysed. 

The  process  of  carrying  out  a  volumetric  estimation  by  the  use 
of  standard  solutions,  is  called  titration ;  thus,  a  solution  of  a 
chloride  is  titrated  with  standard  silver  nitrate  ;  a  ferrous  salt  in 
solution  is  titrated  with  potassium  permanganate  ;  and  so  on. 
Standard  solutions  themselves  have  to  be  titrated,  i.e.  their  chemical 
value  or  power  has  to  be  determined  in  order  to  ascertain  their  exact 
strength.  They  are,  therefore,  often  spoken  of  as  titrated  solutions. 

Classification  of  Volumetric  Methods.— Almost  all  volu- 
metric processes  are  based  upon  one  of  three  principles,  and  they 
may,  therefore,  be  conveniently  classified  in  the  following  manner: — * 

*  One  important  process  which  is  not  included  in  this  classification,  is  the 
estimation  of  copper  by  means  of  potassium  cyanide.  A  description  of  this 
method  will  be  found  on  p.  374. 


302  Volumetric  Methods. 

I.  Methods  based  upon  the  neutralisation  of  alkalies  and  acids, 
known  also  as  methods  of  saturation.     These  methods  are  some- 
times subdivided  into  two  classes,  which  are  included  under  the 
terms  alkalimetry  and  acidimetry,  the  one  being  the  converse  of 
the  other. 

II.  Methods  based  upon  processes  of  oxidation  or  reduction.  . 

III.  Methods  based  upon  precipitation,  either  alone  or  in  con- 
junction with  one  of  the  other  methods.     For  example,  silver  may 
be  estimated  by  precipitation  alone  by  means  of  standard  sodium 
chloride.    As  the  precipitate  readily  settles,  the  point  of  completion 
of  the  reaction  may  be  determined  by  the  failure  of  a  drop  of  the 
standard  reagent  to  produce  any  further  turbidity.     Barium,  on  the 
other  hand,  may  be  determined  by  precipitation  as  carbonate  (from 
a  neutral  solution)  by  the  use  of  an  excess  of  a  standard  solution 
of  sodium  carbonate.     The  precipitate  is  filtered  off,  and  the  excess 
of  sodium  carbonate  in  the  filtrate  and  washing  is  estimated  by 
titration  with  a  standard  acid.     The  excess  so  found,  deducted  from 
the  total  alkali  used,  gives  the  amount  required  for  the  precipitation 
of  the  barium.    Or  the  precipitated  barium  carbonate,  after  thorough 
washing,  is  dissolved  in  an  excess  of  standard  hydrochloric  acid, 
and  the  excess  of  acid  estimated  by  means  of  standard  alkali.     The 
amount  of  acid  required  to  dissolve  the  barium  carbonate  is  thus 
found,  and  by  calculation  the  weight  of  barium  is  ascertained. 

Direct  and  Indirect  Determinations.— In  a  large  number 
of  volumetric  analyses,  substances  are  determined  by  indirect  pro- 
cesses ;  they  are  estimated,  as  it  were,  by  proxy.  The  difference 
between  direct  and  indirect  processes  will  be  most  readily  under- 
stood from  the  following  illustrations. 

Ammonia  may  be  directly  determined  in  an  ammoniacal  salt 
by  heating  a  known  quantity  of  the  salt  along  with  caustic  alkali, 
and  causing  the  whole  of  the  liberated  ammonia  to  pass  into,  and 
be  absorbed  by,  a  measured  volume  of  standard  sulphuric  acid. 
The  exact  quantity  of  this  acid,  which  is  saturated  or  neutralised 
by  the  ammonia,  is  then  ascertained  by  titrating  the  solution  with 
a  standard  alkali.  From  the  amount  of  alkali  thus  used,  the  excess 
of  sulphuric  acid  present  is  ascertained,  and  this  deducted  from  the 
total  volume  of  acid  gives  the  quantity  which  has  been  neutralised 
by  the  ammonia  ;  and  as  the  weight  of  ammonia  to  which  each 
cubic  centimetre  of  the  standard  acid  is  equivalent  is  known,  it  also 
gives  directly  the  weight  of  ammonia  present  in  the  ammoniacal 
salt  taken  for  the  estimation. 

Indirectly  ammonia  can  be  determined  by   boiling  a  known 


Preliminary  Manipulations.  303 

quantity  of  the  ammonium  salt  with  a  measured  volume  of  standard 
caustic  alkali,  and  allowing  the  ammonia  to  escape.  When  the 
whole  of  the  ammonia  has  thus  been  expelled,  the  excess  of  caustic 
alkali  is  estimated  by  titration  with  standard  acid.  The  amount  of 
alkali  thus  determined,  deducted  from  the  original  volume  employed, 
represents  the  quantity  of  alkali  which  has  been  decomposed  and 
converted  into  a  salt  of  the  acid  which  was  originally  in  combination 
with  the  ammonia,  e.g.  if  ammonium  chloride  be  the  salt  employed, 
then  a  certain  quantity  of  the  standard  caustic  soda  will  be  con- 
verted into  sodium  chloride,  and  this  gives  indirectly  the  amount 
of  ammonia. 

In  the  direct  process  we  actually  estimate  the  ammonia,  while 
by  the  indirect  method  we,  in  reality,  determine  the  quantity  of  the 
acid  radical  with  which  the  ammonia  is  combined,  and  from  this 
find  the  ammonia  by  calculation. 

Again,  chromium  in  a  chromate  may  be  determined  indirectly 
by  boiling  the  salt  with  hydrochloric  acid,  whereby  chlorine  is 
liberated  in  the  proportion  of  3  equivalents  of  Cl  for  every  i  equiva- 
lent of  chromium  trioxide,  CrO3  ;  thus — 

K2CrO4  +  8HC1  =  2KC1  +  CrG3  +  4H2O  +  3C1 

The  liberated  chlorine  is  made  to  pass  into  potassium  iodide 
solution,  whereby  an  equivalent  quantity  of  iodine  is  liberated,  and 
the  amount  of  iodine  so  set  free  is  determined  by  means  of  a 
standard  solution  of  sodium  thiosulphate  (as  explained  later). 

In  this  process,  iodine  is  actually  determined,  and  from  this 
indirectly  the  chlorine,  and  yet  more  indirectly  the  chromium  is 
estimated. 

In  order  to  carry  out  volumetric  analyses,  the  three  following 
conditions  must  be  fulfilled,  namely  :  (i)  the  means  of  accurately 
measuring  the  volumes  of  liquids  ;  (2)  the  means  of  preparing  the 
necessary  standard  solutions  and  of  testing  their  accuracy  ;  and  (3) 
the  means  of  readily  ascertaining  the  exact  point  when  the  various 
chemical  reactions  involved  are  complete. 

I.  Instruments  for  Measuring  Liquids.— Graduated 
glass  vessels  are  employed  for  measuring  the  volumes  of  liquids, 
four  different  forms  of  apparatus  being  in  common  use  for  different 
purposes,  namely,  flasks,  cylinders,  pipettes,  and  burettes. 

(a)  Graduated  Flasks.  — These  are  flasks  of  such  a  size  that  they 
shall  contain  a  specified  volume  of  liquid  when  filled  up  to  a  gradu- 
ation mark  upon  the  neck.  They  should  be  provided  with  a 
stopper,  the  neck  should  be  somewhat  long  and  narrow,  and  the 


304 


Volumetric  Methods. 


FIG.  64. 


graduation  mark  should  lie  fairly  low  down  upon  the  neck,  in  order 
that  the  contents  of  the  flask,  when  it  is  filled  to  the  mark,  may  be 
conveniently  shaken  up  (Fig.  64).  The  three  sizes  most  convenient 
for  ordinary  work  are  the  litre,  half- litre,  and 
quarter-litre,  besides  which  a  loo-c.c.  flask  is 
useful. 

If  one  of  these  flasks  be  filled  with  water  up 
to  the  graduation  mark,  it  will  be  obvious  that 
when  the  liquid  is  poured  out,  a  certain  small 
proportion  of  it  remains  behind  adhering  to  the 
inside  of  the  flask ;  in  other  words,  the  flask  does 
not  deliver  absolutely  the  whole  of  the  liquid  it 
contained,  hence,  when  such  a  flask  is  employed 
to  deliver  a  volume  exactly  equal  to  its  indicated 
capacity,  it  must  contain  an  excess  of  the  liquid 
equal  to  the  quantity  which  remains  adhering  to 
the  glass.  Measuring-flasks,  therefore,  usually 
contain  two  graduation  marks,  the  lower  one  being 
the  mark  to  which  the  vessel  must  be  filled  in 
order  to  contain  the  volume  represented  by  the 
denomination  of  the  flask,  while  the  one  slightly  above  it  indi- 
cates the  point  to  which  the  flask  must  be  filled  in  order  that 
it  shall  deliver  that  volume.  This  applies  more  especially  to  the 
smaller  sizes,  as  the  larger  vessels  are  more  exclusively  used -for 
making  up  solutions  to  a  given  volume.  Measuring-flasks  are 
usually  graduated  at  15°  or  15-5°  C.  (60°  F.).  Thus,  a  litre  flask  is 
graduated  to  contain  1000  c.c.  of  water  measured  at  a  temperature 
of  1 5 '5°,  or  about  the  average  temperature  of  the  air.*  Since,  how- 
ever, liquids  undergo  an  appreciable  alteration  in  volume  by  change 
of  temperature,  their  measurements  should  always  be  made  as 
nearly  as  possible  at  one  uniform  temperature,  otherwise  it  is 

*  This  volume  of  water  will  obviously  weigh  slightly  less  than  1000  grams, 
since  i  gram  is  the  weight  of  i  c.c.  at  its  point  of  maximum  density,  namely, 
4°.  Its  actual  weight  is  found  by  multiplying  the  number  of  cubic  centimetres 
by  the  density  of  water  at  15*5°  (given  in  the  table  in  the  Appendix) ;  thus,  1000  x 
0-99911  =  999*11  grams.  Similarly,  for  a  ^-litre  flask,  250x0-99911  =  24977 
grams. 

If  icoo  grams  of  water  at  a  temperature  of  15-5°  be  placed  into  a  flask,  and 
the  volume  it  occupies  be  called  icoo  c.c.,  obviously  these  cubic  centimetres  are 
not  the  true,  or,  as  they  are  termed,  the  absolute  cubic  centimetres  ;  they  are, 
of  course,  slightly  greater— greater  in  the  proportion  of  i  :  1-00089  (this  being 
the  coefficient  of  expansion  of  water  at  15-5°).  In  this  case,  the  so-called  litre 
flask  would  in  reality  have  a  capacity  of  1000-89  absolute  cubic  centimetres ; 
and  similarly,  the  J-litre  flask,  while  containing  150  of  the  larger  arbitrary  cubic 
centimetres,  would  hold  250  X  1-00089  =  250-22  absolute  cubic  centimetres. 


Calibration  of  Litre  Flasks.  305 

necessary  to  introduce  a  correction  for  temperature.  For  example, 
suppose  a  250  c.c.  flask,  which  has  been  graduated  at  15*5°,  is 
filled  to  the  mark  with  water  having  a  temperature  of  20°,  then  from 
the  table  in  the  Appendix,  giving  the  coefficients  of  expansion  of 
water  (the  volume  at  4°  =  i),  we  get  the  following  proportion  : — 

1-00169  :  1-00089  :  :  250  :  249-8  c.c. 

(vol.  at  20°)         (vol.  at  is'5°) 

hence  the  volume  of  liquid  measured  at  the  higher  temperature  is 
about  \  of  a  cubic  centimetre  short.  For  all  ordinary  analytical 
purposes,  however,  slight  variations  of  temperature  to  the  extent  of 
one  or  two  degrees  on  either  side  of  I5'5°  may  be  disregarded. 

The  accuracy  of  the  graduations  of  measuring-flasks  should  be 
tested,  i.e.  the  vessels  should  be  calibrated  before  being  used.  For 
this  purpose  the  dry  clean  flask  is  first  counterpoised  upon  a  balance 
capable  of  carrying  a  heavy  load.  Weights  are  then  placed  upon 
the  scale-pan,  along  with  the  tare,  equal  to  the  number  of  grams  of 
water  at  15-5°  which  the  flask  is  intended  to  contain.*  Water  is 
then  poured  into  the  flask  (which  should  be  removed  from  the 
balance  while  being  filled)  until  the  level  of  the  lowest  point  of 
the  meniscus  (see  p.  310)  is  coincident  with  the  graduation  mark. 
If  the  flask  so  filled  exactly  balances  the  weights,  the  graduation 
may  be  taken  as  correct ;  if  not,  water  is  either  withdrawn  or  added 
by  means  of  a  fine  pipette  until  equilibrium  is  established,  and  a 
fresh  graduation  is  made  upon  the  neck  by  means  of  a  scratching 
diamond  or  a  file. 

In  order  to  graduate  a  flask  to  deliver  a  definite  volume,  exactly 
the  same  procedure  is  carried  out,  except  that  the  flask,  instead  of 
being  dry  at  the  beginning,  is  first  filled  with  water  to  the  mark, 
and  then  emptied  out  and  allowed  to  drain  for  a  few  moments  (as 

*  In  graduating  or  calibrating  measuring-flasks,  one  of  two  plans  may  be 
adopted — 

(1)  The  number  of  grams  of  water  equal  to  the  denomination  of  the  flask 
at  15 '5°  (or  any  other  temperature  which  would  be  more  suitable  under  special 
climatic  conditions)  are  carefully  weighed   out,  and   the  volume  which   they 
occupy  in  the  flask  is  indicated  by  a  mark  scratched  upon  the  neck.     The  cubic 
centimetres  in  this  case  have  an  arbitrary  value,  as  explained  in  the  footnote  on 
p.  304,  but  if  all  the  measuring  vessels  are  graduated  on  the  same  system,  so 
that  this  arbitrary  unit  has  the  same  value  in  them  all,  no  error  will  arise  on 
this  account. 

(2)  A  quantity  of  water  at  15 '5°  is  weighed  out,  which,  if  cooled  to  4°,  would 
then  occupy  a  volume  equal  to  the  denomination  of  the  flask.     The  number  of 
grams  to  be  weighed  out  is  ascertained  by  multiplying  the  denomination  of  the 
flask  by  the  density  of  water  at  15*5°,  as  explained  in  a  previous  footnote. 
In  this  case,  the  cubic  centimetres  represent  the  true  or  absolute  cubic  centi- 
metres.    The  first  method  of  graduating  is  the  one  most  usually  adopted. 


306         Preliminary  Volumetric  Manipulations. 

long  as  it  would  be  allowed  to  drain  when  being  actually  used  to 
deliver).  The  flask  is  then  counterpoised  with  the  remaining  traces 
of  water  adhering  to  the  interior  surface  of  the  glass. 

(b]  Graduated  Cylinders. — These  are  sometimes  used  instead 
of  flasks,  when  less  accurate  measurement  is  all  that  is  necessary. 
Two  forms  are  in  common  use,  as  shown  in  Fig.  65,  the  larger 
stoppered  vessel  being  known  as  a  test-mixer.  As  shown  in  the 
figure,  these  cylinders  are  graduated  throughout  their  length  into 
a  number  of  small  subdivisions  of  the  total  capacity.  The  correct- 
ness of  these  graduations  may  be  tested  by  introducing  successive 


iftetf 

\15'.C\ 


FIG.  65. 


FIG.  66. 


small  volumes  of  water  from  a  pipette  or  a  burette,  which  has  been 
previously  tested.  If  the  graduations  upon  the  cylinder  do  not 
coincide  with  the  volumes  which  are  thus  introduced,  a  table  must 
be  drawn  up  giving  the  actual  value  of  each  graduation. 

(c]  Pipettes. — The  pipette  is  a  glass  tube,  usually  with  a  bulb 
or  enlargement  upon  it,  drawn  to  a  point  at  one  end.  Fig.  66 
shows  the  most  usual  forms  of  a  small  and  a  large  pipette.  The 
instrument  is  graduated  to  deliver  its  specified  volume  of  liquid. 


Calibration  of  Pipettes.  307 

It  is  filled  by  sucking  the  liquid  up  until  it  is  somewhat  higher  than 
the  graduation  mark  upon  the  stem,  and  covering  the  upper  end 
with  the  finger,  as  shown  in  the  figure.  Then,  by  a  slight  release 
of  the  pressure  of  the  finger,  the  liquid  is  allowed  to  flow  slowly  out, 
until  the  level  is  coincident  with  the  mark  upon  the  stem. 

When  the  pipette  is  allowed  to  empty  itself,  the  last  drops 
which  remain  in  the  pointed  end  may  either  be  blown  out,  or  made 
to  flow  out  by  drawing  the  point  along  against  the  wet  sides  of  the 
vessel  into  which  the  liquid  is  being  delivered.  The  former  plan  is 
the  quicker,  and  is  the  one  which  is  most  instinctively  resorted  to  ; 
but  the  latter  is  the  more  accurate  method.  It  is  not  a  matter  of 
serious  importance  which  method  is  adopted,  so  long  as  the  same 
plan  is  uniformly  employed. 

Pipettes  are  calibrated  to  deliver  a  definite  volume  by  a  method 
practically  the  same  as  in  the  case  of  flasks.  To  carry  it  out,  the 
pipette  is  first  filled  with  water  by  suction,  and  the  end  of  the  tube 
which  was  dipped  into  the  water  is  wiped  dry  on  the  outside.  The 
liquid  is  then  allowed  to  run  out,  and  the  last  drops  removed  by 
one  of  the  above-mentioned  plans.  The  point  of  the  instrument  is 
now  closed  by  placing  a  minute  fragment  of  wax  *  upon  it,  and 
then  just  softening  the  wax  by  bringing  the  point  near  to  a  small 
flame.  In  this  way  the  pipette  can  be  closed  without  the  wax 
entering  the  tube.  The  moist  pipette  is  then  placed  upon  a  balance 
and  counterpoised. 

It  is  next  filled  with  water  at  a  temperature  of  I5'5°,  which  can 
be  introduced  by  means  of  a  glass  tube  drawn  out  sufficiently  fine 
to  pass  down  the  stem  of  the  pipette.  After  the  first  few  drops 
have  been  put  in,  the  air-bubble  which  is  usually  trapped  in  the 
bottom  of  the  tube  may  generally  be  dislodged  by  a  few  shakes  ;  if 
this  fails,  however,  it  may  be  removed  and  the  water  made  to  run 
right  into  the  point  by  thrusting  a  long  capillary  tube  down  to  the 
bottom.  Care  must  be  taken,  however,  that  no  fragments  of  this 
fine  tube  become  broken  off  and  left  in  the  apparatus.  During  the 
filling  process  the  pipette  should  be  supported  by  a  clamp,  and  not 
held  in  the  warm  hand,  otherwise  the  temperature  of  the  water 
may  be  considerably  raised  during  the  operation.  When  the  water 
is  coincident  with  the  graduation  mark,  the  instrument  is  carefully 
laid  upon  the  scale-pan  (the  water  will  not  run  out  of  it,  although 
placed  in  a  horizontal  position),  and  gram  weights  equal  to  the 
number  of  cubic  centimetres  the  pipette  is  intended  to  deliver  are 

*  The  black  bituminous  material  known  as   "bicycle  cement"  answers 
admirably  for  this  purpose. 


308          Preliminary  Volumetric  Manipulations. 


placed  upon  the  opposite  pan.  If  this  is  found  to  exactly  counter- 
poise the  apparatus,  the  graduation  may  be  taken  as  correct ;  but 
if  not,  water  must  be  either  withdrawn  or  added,  as  the  case  may 
be,  until  the  apparatus  is  exactly  equipoised,  and  a  fresh  mark 
scratched  upon  the  stem. 

(d]  Burettes. — The  burette  is  a  long  straight  glass  tube,  one 
end  of  which  is  drawn  down  and  terminated  by  a  glass  stop-cock, 
or  connected  to  a  jet  by  means  of  a  caoutchouc  tube  which  can  be 

closed  by  means  of  a  pinch- 
cock.  Fig.  67  shows  the  two 
forms  of  apparatus.  The 
burette  is  graduated  almost 
throughout  the  entire  length, 
the  graduations  being  usually 
tenths  of  a  cubic  centimetre, 
as  shown  in  Fig.  68.  The 
size  most  commonly  used 
has  a  capacity  of  50  c.c.  For 
ordinary  use  the  common  re- 
tort-stand and  clamp  shown 
to  the  left  in  the  figure  make 
a  very  convenient  stand  for 
holding  burettes.  If  desired, 
several  can  be  supported 
upon  the  same  retort-stand. 
The  instrument  is  filled  by 
means  of  a  small  funnel 
placed  in  the  top  (which 
should  be  removed  after- 
wards, lest  any  adhering 
drops  fall  into  the  burette), 


'iiiiiiiiiniiiiiiiiiimiii 


FIG.  67. 


until  the  liquid  is  consider- 
ably above  the  topmost  gra- 
duation. The  tap  or  pinch-cock  is  then  momentarily  opened,  in 
order  that  the  liquid  may  sweep  out  before  it  the  air  which  is 
contained  in  the  tap  and  jet.  This  will  not  be  successfully  accom- 
plished if  the  tap  or  pinch-cock  is  only  gradually  opened,  as  the 
liquid  then  slips  down  the  walls  of  the  narrowed  portions  and 
leaves  air-bubbles  in  the  tubes,  which  it  is  absolutely  necessary 
to  remove  before  the  instrument  can  be  used  with  exactness. 

The  calibration  of  a  burette  is  a  more  serious  operation  than 
that  of  a  pipette,  for  since  it  is  obviously  out  of  the  question  to 


Calibration  of  Btirettes.  309 

regraduate  the  instrument,  the  true  value  of  the  graduations  upon 
it  must  be  ascertained,  and  a  special  table  constructed  for  each 
instrument.  The  process  is  carried  out  in  the  following  way  :  The 
burette  is  filled  with  water  at  a  temperature  of  15*5°  (if  possible  in 
a  room  having  that  temperature,  since  the  operation  occupies  a 
considerable  time),  the  air-bubbles  being  swept  out  as  described, 
and  the  level  of  the  liquid  being  coincident  with  the  topmost 
graduation.  A  clean,  dry,  stoppered  flask,  capable  of  containing 
as  much  liquid  as  the  burette,  is  accurately  counterpoised,  and  the 
water  in  the  burette  is  delivered  into  it  in  successive  small  quanti- 
ties of  3,  4,  or  5  of  the  c.c.  graduations  at  a  time.  After  each  small 
portion  of  water  has  been  run  into  it,  the  flask  is  weighed  until  the 
whole  50  c.c.  has  been  delivered. 

The  number  of  grams  which  each  portion  weighs,  represents  the 
number  of  cubic  centimetres  it  actually  occupies,  and  if  this  is 
different  from  the  number  of  c.c.  graduations,  the  error  is  equally 
divided  between  them.*  Thus,  suppose  the  water  is  delivered  in  por- 
tions represented  by  five  of  the  c.c.  divisions,  and  the  first  portion 
is  found  to  weigh  5*15  grams,  instead  of  exactly  5  grams  ;  then  the 
error,  namely  0*15,  is  divided  equally  among  the  five  graduations, 
each  one  of  which  will  then  have  the  value,  not  of  i  c.c.,  but 
1-03  c.c.  The  table,  therefore,  will  run  as  follows  for  the  first  five 
graduations  — 

Graduation  I  =  1*03  c.c. 

2  =  2'06     „ 

»         3  =  3'09    „ 
„         4  =  4-12    „ 

»         5  =  5'iS    » 

Suppose,  after  the  second  portion  has  been  delivered,  the  weight 
of  the  water  is  io'i7  grams  ;  then  — 

1  0*17  —  5*15  =  5*02  grams  =  the  weight  of  the  second  portion  of  water, 


and    -     =  i  '004  c.c.  =  the  value  of  each  graduation  between  5 

and  10 
hence  the  table  will  continue  — 

*  It  will  be  evident  that  the  larger  the  volume  which  is  delivered  at  one 
time,  the  less  accurate  is  the  calibration.  Thus,  from  the  above  illustration, 
if  instead  of  delivering  the  contents  of  five  graduations,  that  of  ten  had  been 
weighed  at  once,  the  value  of  each  graduation  would  have  been— 

^Z=roi7 

10 

and  the  value  of  the  fifth  graduation  would  have  come  out— 
1-017x5  —  5*085  c.c.  instead  of  5^15 


3io         Preliminary  Volumetric  Manipulations. 


Graduation  6  =  5'i 

5)  7    =        55 

55  8    =       „ 

9=    » 

10  =    „ 


•004  =  6' 1 54  c.c. 
•008  =    7-158   „ 

•012  =  8*162    ,, 

•016  =  9*166    , 

•020  =  io' 1 70  „ 


Small  burettes  holding  10  c.c.  are  sometimes  used  after  the 
manner  of  a  pipette.  They  consist  of  straight  tubes  drawn  to  a 
point  at  one  end,  and  graduated  into  tenths  of  a  cubic  centimetre, 
as  in  the  burette.  They  are  filled  by  suction,  and  employed  as 
pipettes  for  delivering  various  small  measured  volumes. 

Beading  the  Volume  of  Liquids  in  Graduated  Vessels. 
— Owing  to  the  action  of  capillarity,  the  surface  of  a  liquid  con- 
tained in  a  glass-tube  is  not  plane,  but  curved,  the  extent  to  which 
it  is  curved  depending  (for  the  same  liquid)  upon  the  diameter  of 
the  tube. 

This  curved  surface  is  called  the  meniscus.  Figs.  68  and  70 
show  the  meniscus  in  the  case  of  water  contained  in  a  burette. 


15-3 


16- 


18- 


FIG.  68. 


FIG.  69. 


Ill  reading  graduated  instruments,  it  is  usual  to  take  the  gradu- 
ation which  coincides  with  the  lowest  point  of  the  meniscus.  Thus, 
in  Figs.  68  and  70  the  volume  of  the  liquid  would  be  taken  as  16*5. 
If,  as  sometimes  happens  in  certain  lights,  the  line  of  the  curve  is 
not  very  clearly  visible,  it  may  usually  be  made  quite  distinct  by 


The  Use  of  Floats.  3 1 1 


holding  behind  the  tube  a  piece  of  white  paper,  inclining  it  slightly 
upwards,  as  in  Fig.  68.  In  order  to  make  a  correct  reading,  it  is 
necessary  that  the  eye  of  the  observer  should  be  in  the  same  hori- 
zontal plane  as  the  surface  of  the  liquid.  The  reason  for  this  will 
be  evident  from  the  diagram  (Fig.  69),  where  an  error  of  one  gradu- 
ation would  arise  by  reading  from  the  positions  A  or  A1  instead  of 
from  B. 

Where  the  graduation  mark  is  continued  as  a  ring  right  round 
the  vessel,  as  is  usually  the  case  with  pipettes  and  flasks,  the  eye 
will  be  situated  in  a  horizontal  plane  with  the  graduation,  when  the 
front  of  the  mark  exactly  overlays  the  back,  so  that  the  ring 
appears  as  a  line.  The  lowest  point  of  the  meniscus  must  then 
just  touch  this  line.  A  simple  device  to  ensure  that  the  readings 
of  a  burette  (where  the  graduations  are  on  one  side  of  the  tube 
only)  shall  be  consistently  made  from  a  correct  point  of  observa- 
tion is  shown  in  Fig.  70.  It  consists  merely  of  a  narrow  strip  of 
card  folded  in  the  middle,  with  the  two  free  ends  pinned  together 
with  a  paper-fastener.  This  is  slipped  over  the  burette,  and  while 
it  allows  of  being  easily  slid  up  and  down  the  tube,  it  will  also  hold 
itself  in  any  position  in  which  it  may  be  put.  When  taking  a  read- 
ing, this  little  clip  is  placed  so  that  its  upper  edge  is  just  a  little 
below  the  level  of  the  liquid  :  then,  when  the  eye  is  in  such  a 
position  that  the  back  and  front  edges  of  the  clip  just  coincide,  it 
will  also  be  practically  in  the  same  horizontal  plane  as  the  bottom 
of  the  meniscus. 

The  same  object  may  be  gained  by  the  use  of  floats.  These  are 
little  weighted  glass  bulbs  or  tubes  which  are  placed  inside  the 
burette,  floating  upon  the  liquid.  The  graduation  which  coincides 
with  the  horizontal  ring-mark  upon  the  float  is  the  one  which  is 
read.  Two  forms  of  floats  are  seen  in  Fig.  71.  A  represents  the 
most  familiar  shape.  Unless  some  care  is  exercised,  the  use  of 
such  a  float  is  very  liable  to  introduce  errors.  If  it  is  much 
narrower  than  the  burette,  it  may  sometimes  take  up  positions  in 
which  the  graduation  mark  upon  it  is  not  perfectly  horizontal.  On 
the  other  hand,  if  it  is  too  close  a  fit,  it  is  very  prone  to  lag  behind 
the  liquid,  or  even  to  stick  altogether.  B  (Fig.  71)  represents  a 
newer  and  better  form  of  float.  Owing  to  its  shape  it  always  main- 
tains a  vertical  position,  and  may  therefore  be  sufficiently  narrow 
at  its  widest  point  to  admit  of  perfect  freedom  of  movement  in  the 
burette.  The  graduation  mark  on  this  float  is  upon  the  smaller 
bulb  at  the  top,  which  projects  above  the  liquid  altogether,  hence 
this  form  of  apparatus  may  be  used  equally  well  whether  the  liquid 


312          Preliminary  Volumetric  Manipulations. 


in  the  burette  is  colourless  or  even  opaque.     The  reading,  as  the 
float  stands  in  the  figure,  is  14*4. 

It  is  very  necessary,  whatever  float  is  used,  to  keep  a  close 
watch  upon  it  as  the  liquid  in  the  burette  gets  down  towards  the 
bottom  graduations.  It  is  well  to  place  an  indiarubber  ring  or 
other  convenient  mark  upon  the  burette  at  the  point  beyond  which 
it  is  not  admissible  to  go  when  the  float  is  being  used. 


FIG.  70. 


FIG.  71. 


II.  Standard  Solutions.— Solutions  of  known  strength  which 
are  used  in  volumetric  analysis  are  called  standard  solutions. 
These  various  solutions  are  usually  made  of  such  strengths  that 
the  quantities  of  the  positive  or  negative  constituents,  or  of  what 
may  be  called  the  active  constituents,  of  the  compounds  in  the  solu- 
tions shall  bear  the  same  relation  to  each  other  as  the  numbers 
which  express  their  chemical  equivalence — that  is  to  say,  equal 
volumes  of  the  different  solutions  will  contain  eqitivalent  propor- 
tions of  the  effective  substances  in  solution. 

For  example,  standard  solutions  of  sodium  chloride  and  silver 
nitrate  will  be  of  such  strengths  that  whatever  volume  of  the  former 
contains  35*5  grams  of  chlorine,  the  same  volume  of  the  other 
shall  contain  108  grams  of  silver  ;  or,  again,  standard  solutions  of 


Standard  Solutions.  313 

hydrochloric  acid,  of  sodium  hydroxide,  and  of  sodium  carbonate 
will  be  of  such  strengths  that  whatever  volume  of  the  first  contains 
i  gram  of  hydrogen,  the  same  volume  of  the  others  shall  each 
contain  23  grams  of  sodium. 

Normal  Standard  Solutions. — When  the  solutions  are  of 
such  a  strength  that  the  particular  volume  which  contains  I  gram 
of  hydrogen,  23  grams  of  sodium,  35-5  grams  of  chlorine,  etc.,  is 
one  litre,  then  the  solution  is  known  as  a  normal  solution.  Thus,  a 
normal  solution  of  hydrochloric  acid  will  contain  I  gram  of  hydro- 

H        ci 

gen  in  the  litre,  therefore  it  must  contain  i  +  35-5  =  36*5  grams  of 
HC1  in  that  volume.  Again,  normal  caustic  soda  will  contain 

Na        H          O 

23  grams  of  sodium  per  litre,  therefore  it  must  contain  23  +  1  +  16 
=  40  grams  of  NaHO  in  the  1000  c.c.  Or,  again,  normal  sodium 
carbonate  must  also  contain  23  grams  of  sodium  per  litre  ;  there- 

Na2        C         O3 

fore  the  volume  must  contain  4-     '——^     -  =  53  grams  of  Na2CO3. 

Similarly,  if  normal  sulphuric  acid  is  to  be  of  such  a  strength 
that  I  litre  shall  contain  i  gram  of  hydrogen,  or,  in  other  words, 
that  the  amount  of  the  negative  radical  SO4  shall  be  the  chemical 
equivalent  of  i  gram  of  hydrogen,  then  the  litre  must  contain 
H3  s  o, 

2  +  32  +  64  =  49  grams  of  H2S04< 

In  some  cases,  it  is  only  a  portion,  and  not  the  whole,  of  some 
particular  element  or  radical  present  in  the  solution,  which  takes 
an  active  part  in  the  chemical  reactions  for  which  the  solution 
is  used.  Thus,  in  the  case  of  potassium  dichromate,  K2Cr2O7,  only 
three  out  of  the  seven  oxygen  atoms  contained  in  the  salt  are 
available  for  purposes  of  oxidation.  We  may  regard  the  compound 
as  breaking  up  into  K2O,Cr2O3,3O.  A  normal  solution,  therefore, 
of  this  salt  (that  is,  one  which  shall  contain  per  litre  a  weight  of 
available  oxygen  chemically  equivalent  to  i  gram  of  hydrogen) 

Ka  Cra  07 

.      78  +   I04'8  +  112  -     .  .      . 

must  contain  —  ^  -  =  49*13  grams  of  the  salt  m  every 

loco  c.c.  The  formula-weight  of  the  salt,  in  this  instance,  being 
divided  by  6  instead  of  by  3,  because  the  equivalent  of  oxygen  is  8, 
and  not  16. 


Preliminary  Volumetric  Manipulations. 

Again,  the  available  oxygen  in  potassium  permanganate  is  five 
out  of  eight  atoms,  2(KMnO4)  =  K2O,2MnO,5O.  Hence  a  normal 
solution  of  this  salt,  containing  in  the  litre  a  weight  of  available 
oxygen  equivalent  to  I  gram  of  hydrogen  (i.e.  8  grams  of  available 
oxygen),  must  contain  in  that  volume  of  the  solution  a  weight  of 
salt  in  grams  equal  to  one-tenth  of  the  formula- weight  K2Mn2O8,  or 

one-fifth  of  that  of  KMnO4,  namely,  —  =•  31-6  grams.* 

Solutions  of  one-half,  one-tenth,  and  one-hundredth  of  the 
strength  of  a  normal  solution  are  called  respectively  semi-normal, 
deci-normal,  and  centi-normal  solutions,  and  are  distinguished  by 

N    N         .    N. 
the  signatures  — '  — '  and  — 

2      10  100 

In  preparing  standard  solutions,  the  weighed  quantity  of  salt  is 
not  added  to  and  dissolved  in  the  measured  volume  of  water,  but 
it  is  first  dissolved  in  a  moderate  quantity  of  water,  and  the  solution 
afterwards  diluted  up  to  the  requisite  volume  by  the  addition  of 
water.  At  first  sight  it  might  be  supposed  that  the  two  processes 
would  give  the  same  result,  but  this  is  not  the  case,  because  the  act 
of  solution  of  many  substances  is  attended  by  a  change  (usually  a 
contraction)  in  the  volume,  and  therefore  in  one  case  the  solution 
obtained  would  be  different  in  volume,  and  therefore  in  strength, 
from  that  prepared  in  the  other  way. 

Further  details  for  the  preparation  of  the  various  standard  solu- 
tions will  be  given  in  the  sections  devoted  to  the  special  analytical 
processes  for  which  the  solutions  are  employed. 

(III.)  Methods  for  ascertaining  the  Completion  of 
Volumetric  Reactions. — It  is  essential  to  the  success  of  a  volu- 
metric process  that  the  exact  point  when  the  chemical  reaction  is 
complete  should  be  readily  determinable.  This  is  usually  accom- 
plished by  means  of  a  third  substance,  known  as  an  indicator,  which 
is  effected  in  some  visible  way  either  by  the  volumetric  reagent  or 
by  the  solution  of  the  substance  being  estimated.  For  example,  in 
order  to  determine  the  exact  point  when  an  alkali  has  been  neu- 
tralised by  an  acid,  a  small  quantity  of  litmus  solution  may  be  added 

*  Some  chemists,  unfortunately,  apply  the  term  normal  to  solutions  which 
contain  in  one  litre  the  formula-weight  in  grams  (i.e.  the  gram-molecule)  of  the 
compound.  According  to  this  method,  while  a  normal  solution  of  caustic  soda 
contains  23  grams  of  sodium  per  litre,  the  normal  sodium  carbonate  will  contain 
46  grams  of  sodium  in  the  same  volume.  To  prevent  confusion,  these  solutions 
should  be  distinguished  from  the  true  normal  solutions  by  the  name  molecular 
solutions. 


The  Use  of  Indicators.  315 

to  the  solution.  Again,  the  point  of  complete  precipitation  of 
chlorine  by  means  of  standard  silver  nitrate,  can  be  ascertained 
with  exactness  by  adding  a  small  quantity  of  a  solution  of  potassium 
chromate  to  the  solution  containing  the  chloride  to  be  estimated. 
The  silver  combines  by  preference  with  the  chlorine,  forming  silver 
chloride  ;  but  when  the  whole  of  the  chlorine  has  been  precipitated 
as  silver  chloride,  then  the  silver  attacks  the  chromate,  precipitating 
the  red  silver  chromate — hence  the  production  of  a  red  precipitate 
which  persists,  is  the  sign  that  the  precipitation  of  the  chlorine  is 
complete.  The  reagent,  or  standard  solution,  may  itself  be  the 
indicator;  this  is  the  case,  for  example,  with  potassium  perman- 
ganate, in  which  the  characteristic  violet  colour  of  the  solution 
serves  to  show  the  precise  moment  when  the  oxidising  process  is 
complete.  Very  often  an  outside  indicator  is  made  use  of ;  that  is 
to  say,  instead  of  adding  the  indicator  to  the  solution  which  is  being 
tested,  the  progress  of  the  reaction  is  watched  by  removing  a  drop 
of  the  liquid  upon  a  glass  rod  at  frequent  intervals,  and  applying  it 
to  the  indicator.  Thus,  in  estimating  the  chlorine  in  bleaching 
powder  by  the  oxidation  of  sodium  arsenite,  the  indicator  employed 
is  a  mixture  of  potassium  iodide  and  starch,  which  gives  the 
characteristic  blue  colour  when  in  contact  with  chlorine.  The 
iodide,  however,  cannot  be  added  to  the  solution  of  bleaching 
powder,  but  by  bringing  a  drop  of  the  latter  into  contact  with 
potassium  iodide  and  starch  paper  until  it  no  longer  produces  the 
blue  coloration,  it  is  easy  to  hit  the  exact  point  when  the  oxidation 
of  the  arsenz^wj-  salt  into  the  arsen/V  state  is  completed. 

Similarly,  the  point  when  a  ferrous  salt  is  completely  oxidised 
to  the  ferric  condition  by  means  of  potassium  dichromate,  is  ascer- 
tained by  withdrawing  a  drop  of  the  mixture  upon  a  glass  rod  and 
applying  it  to  a  drop  of  potassium  ferricyanide  solution  upon  a 
white  plate  or  porcelain  tile  ;  so  long  as  any  ferrous  salt  is  present, 
a  blue  coloration  is  produced,  but  as  soon  as  the  iron  is  entirely 
oxidised,  a  drop  of  the  solution  so  applied  to  the  ferricyanide 
indicator  will  produce  no  coloration. 


SECTION   II. 

VOLUMETRIC  METHODS  BASED  UPON  SATURATION. 
Alkalimetry  and  Acidimetry. 

Normal  Alkalies  and  Acids. — The  four  following  standard 
solutions  may  be  prepared,  namely,  normal  sodium  carbonate, 
sodium  hydroxide,  sulphuric  acid,  and  hydrochloric  acid. 

In  preparing  such  a  series  of  solutions,  it  is  necessary  to  select 
one  of  them  as  the  foundation  or  basis  for  the  others  ;  and  the  con- 
ditions which  must  be  fulfilled  by  the  one  which  is  chosen  for  this 
purpose  are,  that  it  admits  of  being  made  up  with  the  utmost  possible 
accuracy.  Of  these  four  solutions  the  alkaline  carbonate  is  the  one 
which  most  satisfactorily  fulfils  these  requirements,  for  it  is  a  salt 
which  can  readily  be  obtained  in  a  state  of  practical  purity,  and  it 
also  admits  of  being  weighed  with  perfect  exactness. 

Having  once  obtained  a  normal  solution  of  sodium  carbonate 
of  exact  strength,  it  can  be  employed  as  a  basis  for  the  preparation 
of  normal  sulphuric  and  hydrochloric  acids,  and  by  means  of  either  of 
these  the  standardisation  of  the  caustic  alkali  may  be  effected. 
The  strength  of  the  acids  can  be  checked,  if  desired,  by  precipitation 
(the  one  with  barium  chloride,  and  the  other  with  silver  nitrate) 
and  gravimetric  determination. 

Indicators. — There  are  a  number  of  substances  which  may 
be  employed  in  order  to  indicate  the  exact  point  of  neutrality,  or 
the  end  reaction,  as  it  is  termed,  in  processes  of  alkalimetry  and 
acidimetry.  The  most  important  and  useful  *  are  litmus,  methyl 
orange,  and  phenolphthalein. 

Litmus  Solution. — This  is  prepared  by  crushing  about  10  grams 
of  the  solid  in  a  clean  mortar  with  hot  distilled  water.  The  liquid 
is  poured  off,  and  the  residue  further  extracted  by  one  or  two  fresh 
supplies  of  hot  water.  The  extract  is  then  allowed  to  settle  for 
some  hours,  and  the  clear  blue  liquid,  diluted  to  about  200  c.c.,  is 

*  For  a  full  account  of  the  various  indicators,  their  relative  sensitiveness, 
and  special  uses,  see  Thomson, ,Chem.  News. 


Volumetric  Analysis.  317 

transferred  to  a  bottle.  As  this  solution  loses  its  colour  if  closely 
corked  or  stoppered  up  so  as  to  exclude  the  air,  it  is  a  good  plan  to 
place  in  the  neck  of  the  bottle  a  flat- topped  stopper  which  drops  in 
quite  loosely,  and  serves  merely  as  a  cover  to  prevent  the  entrance 
of  foreign  matter.  A  few  drops  of  chloroform  added  and  shaken 
up  with  the  solution,  will  prevent  mould  from  forming  in  it,  which 
otherwise  the  liquid  is  very  prone  to  develop. 

Litmus  solution  is  turned  red  by  acids,  and  blue  by  alkalies  ; 
the  colour  which  it  should  exhibit  when  perfectly  neutral  is  violet; 
it  is  usually  necessary,  therefore,  to  add  a  drop  or  two  of  very  dilute 
acid  (preferably  nitric  acid)  by  dipping  a  glass  rod  into  the  acid 
and  stirring  it  into  the  litmus  until  a  purple  or  violet  tint  is  imparted 
to  the  solution.* 

When  carbon  dioxide  is  disengaged  (as  when  an  alkaline 
carbonate  is  being  titrated  with  an  acid,  or  vice  versa\  litmus 
cannot  be  employed  as  the  indicator  unless  the  liquid  be  boiled. 
This  is  owing  to  the  fact  that  the  dissolved  carbon  dioxide  itself 
exerts  a  feeble  acid  reaction  upon  the  litmus,  and  it  is  only  when 
the  gas  is  entirely  expelled  by  boiling  that  the  litmus  becomes  the 
true  indicator  of  the  point  of  neutrality.  It  is  under  these  circum- 
stances that  methyl  orange  is  the  more  convenient  indicator. 

Methyl  Orange, — This  substance  dissolves  in  water,  giving 
an  orange-coloured  solution,  which  is  turned  yellow  by  alkalies 
and  a  pink-red  colour  by  acids.  The  solution  for  use  as  an  indi- 
cator is  prepared  by  dissolving  o'5  gram  of  the  solid  in  500  c.c.  of 
water  ;  and  one  or  two  drops  only  should  be  employed,  as  its 
indications  are  less  sensitive  if  much  of  it  be  used. 

Methyl  orange  is  unaffected  by  carbon  dioxide  ;  it  is,  therefore, 
specially  adapted  for  titrations  of  alkaline  carbonates  with  mineral 
acids,  or  vice  versa.  With  this  indicator,  therefore,  these  determi- 
nations can  be  carried  out  in  the  cold. 

Methyl  orange  is  only  suitable  as  an  indicator  where  mineral 
acids  are  concerned  ;  it  cannot  be  used  with  organic  acids. 

Phenolphthalein. — A  solution  of  this  compound  is  prepared 
by  dissolving  i  gram  of  the  solid  in  100  c.c.  of  alcohol.  The  dilute 
neutral  solution  is  colourless,  but  on  the  addition  of  an  alkali  it 

*  The  aqueous  infusion  of  litmus  obtained  in  this  way  contains  other  things 
besides  the  blue  colouring  matter,  and  the  presence  of  these  tends  to  diminish 
the  sensitiveness  of  the  reagent.  The  saline  substances  may  be  removed  by 
acidifying  the  extract  with  hydrochloric  acid,  and  submitting  the  mixture  to 
the  process  of  dialysis.  The  purified  colouring  matter  which  remains  behind 
on  the  dialyser  is  then  dissolved  in  hot  water.  It  is  only  for  very  special 
purposes,  however,  that  it  is  necessary  to  prepare  such  a  purified  litmus 
solution. 


3i8  Volumetric  Methods  of  Analysis. 

becomes  a  deep  red  colour.    This  colour  is  immediately  discharged 
when  the  liquid  is  acidified  either  with  mineral  or  organic  acids. 

The  chief  value  of  this  indicator  lies  in  its  applicability  to  the 
titration  of  organic  acids.  It  cannot  be  employed  in  cases  where 
carbon  dioxide  is  evolved,  since  its  colour  is  destroyed  by  carbonic 
acid  ;  but  as  the  acid  carbonates  (bi-carbonates)  do  not  give  the 
red  colour  with  this  compound,  it  can  be  made  use  of  to  indicate 
the  completion  of  theflrst  stage  in  the  neutralisation  of  a  normal 
carbonate,  namely,  the  conversion  of  the  normal  into  the  acid  car- 
bonate. Phenolphthalein  cannot  be  used  in  the  presence  of 
ammonia. 

Normal  Sodium  Carbonate. 

(53  grams  of  Na2CO3  per  litre.} 

In  order  to  obtain  this  salt  in  a  state  of  purity,  it  is  prepared  by 
heating  the  purest  sodium  bicarbonate  to  a  dull  red  heat  until  no 
further  loss  of  carbon  dioxide  and  water  takes  place.*  Theoretically, 
84  grams  of  bicarbonate  should  yield  53  of  the  normal  salt ;  a 
slight  excess  of  this  proportion,  therefore,  should  be  employed.  To 
prepare  half  a  litre  of  the  standard  solution,  about  43  grams  of  the 
pure  sodium  bicarbonate  are  heated  in  a  weighed  platinum  dish  to 
a  low  red  heat  for  about  10  or  15  minutes.  The  salt  must  not  be 
allowed  to  fuse.  It  is  then  cooled  in  a  desiccator  and  weighed. 
To  ensure  that  the  decomposition  has  been  completed,  the  dish  is 
again  heated  for  another  10  minutes,  and,  after  cooling  in  the 
desiccator,  weighed  again.  The  weight  of  the  salt  (after  deducting 
the  weight  of  the  dish)  will  be  a  little  over  26-5  grams.  By  means 
of  a  clean  spatula  or  pen-knife,  a  small  quantity  is  removed,  so  as 
to  bring  the  weight  to  exactly  26^5  grams,  the  operation  being 
performed  without  undue  exposure  of  the  dry  salt.  The  contents 
of  the  dish  are  then  washed  out  into  a  beaker,  the  dish  being 

*  Owing  to  its  much  slighter  solubility  in  water,  the  bicarbonate  is  the 
more  easily  purified  salt ;  hence  of  the  two  so-called  "  pure  "  salts  of  commerce, 
the  bicarbonate  will  have  a  higher  degree  of  purity  than  the  normal  salt.  The 
sample  employed  here,  however,  must  first  be  tested  in  the  following  way  :  A 
few  grams  are  dissolved  in  a  small  quantity  of  hot  water,  which  at  once  reveals 
the  presence  or  absence  of  any  traces  of  insoluble  impurity.  The  solution  is 
then  acidified  with  nitric  acid  (which  must  be  absolutely  free  from  either  hydro- 
chloric or  sulphuric  acid),  and  tested  for  chlorides  and  sulphates  in  the  usual 
way.  If  the  salt  does  not  dissolve  to  a  perfectly  clear  solution,  and  is  not  free 
from  chlorides  and  sulphates,  a  quantity  of  it  must  be  purified  by  recrystal- 
lisation.  A  hot  saturated  solution  is  made,  and  filtered  by  the  use  of  one  of 
the  arrangements  for  warming  the  funnel  shown  on  p.  210.  The  filtered  liquid 
is  continually  stirred  as  it  cools,  whereby  the  salt  is  deposited  in  a  fine  granular 
condition.  The  mother-liquor  is  then  decanted,  and  the  crystals  drained  and 
dried  upon  a  porous  plate. 


Acidimetry  and  Alkalimetry.  319 

thoroughly  rinsed  with  warm  water,  and  the  salt  completely  dissolved 
by  stirring  the  mixture  with  a  glass  rod.  The  solution  is  then  care- 
fully poured  into  a  half-litre  flask,  the  beaker  being  several  times 
rinsed  with  water.  The  liquid  is  then  cooled  to  15*5°,  and  water 
at  the  same  temperature  is  added,  with  continual  gentle  shaking, 
until  the  neck  is  reached.  The  vessel  is  then  carefully  filled  to  the 
graduation  mark,  after  which  the  stopper  is  inserted,  and  the 
contents  thoroughly  mixed  by  shaking. 

Instead  of  bringing  the  weight  of  the  sodium  carbonate  to  the 
exact  quantity  required  for  half  a  litre  of  solution,  the  whole  of  the 
salt  in  the  dish  may  be  employed,  and  the  exact  volume  of  water 
which  will  be  required  in  order  to  make  the  solution  of  normal 
strength  is  calculated.  Thus,  suppose  the  weight  of  sodium  car- 
bonate in  the  dish  after  heating  is  26749  grams,  instead  of  26*5  ; 
then— 

26'5  :  26749  •  5°°  c-c-  :  5°47  c-c- 

Hence,  after  the  solution  has  been  made  up  to  the  graduation  mark 
in  the  manner  described  above,  47  c.c.  of  water  are  added  from  a 
burette,  and  the  solution  then  finally  shaken  up  to  ensure  thorough 
mixing. 

Normal  Sulphuric  Acid. 
(49  grams  of  HZSO^  per  litre.} 

The  specific  gravity  of  ordinary  oil  of  vitriol  being  about  r8 
49  grams  will  be  rather  less  than  30  c.c.  Hence,  if  this  volume  be 
measured  out  and  diluted  up  to  a  litre,  a  solution  will  be  obtained 
which  will  have  a  rough  approximation  to  the  required  strength. 
If  this  solution  is  then  titrated  with  the  standard  sodium  carbonate, 
its  actual  strength  can  be  ascertained  ;  and  by  calculation,  the 
volume  of  water  which  must  be  added  in  order  to  bring  it  to  the 
exact  normal  strength  is  determined. 

Dilution  of  the  Acid. —  Thirty  cubic  centimetres  of  pure 
sulphuric  acid  are  gradually  poured  into  about  150  c.c.  of  water 
in  a  flask.  The  mixture  is  then  cooled  by  holding  the  flask  under 
the  water-tap,  and  allowing  a  stream  of  water  to  run  over  it,  at 
the  same  time  shaking  the  liquid  round  within  the  vessel.  When 
cold,  the  solution  is  transferred  to  a  litre  flask,*  and  the  volume 
made  up  to  the  1000  c.c.  mark  by  the  addition  of  water  at  a 
temperature  of  I5'5°. 

*  It  is  not  advisable  to  dilute  the  strong  acid  in  the  litre  flask  itself,  as 
there  is  always  some  risk  (although  perhaps  not  much  when  experience  has 
been  gained)  of  fracturing  the  vessel ;  and  it  is  obviously  wiser  that  a  common 
flask,  and  not  one  which  has  been  carefully  graduated,  should  take  this  risk. 


320  Volumetric  Methods  of  Analysis. 

Titratiou  of  the  Acid.— A  burette  is  first  filled  with  the 
dilute  acid,*  care  being  taken  to  remove  all  air-bubbles  from  the 
tap,  as  explained  on  p.  308.  Twenty-five  cubic  centimetres  of  the 
normal  sodium  carbonate  are  then  transferred  by  means  of  a  pipette  * 
to  a  small  beaker,  and  a  single  drop  of  the  methyl  orange  indicator 
(see  p.  317)  is  added.  The  beaker  is  then  placed  upon  a  white 
glazed  tile  beneath  the  burette,  and  the  acid  gradually  run  into  the 
alkaline  liquid,  the  solution  being  gently  rotated  in  order  to  ensure 
thorough  mixing  after  each  addition  of  acid.  At  first  2  or  3  c.c. 
of  the  acid  may  be  added  at  a  time  ;  but  as  the  point  of  neutrality 
is  approached,  smaller  and  smaller  quantities  are  added  at  once, 
until  they  are  reduced  to  a  few  drops  only,  and  at  last  the  addition 
of  a  single  drop  produces  a  permanent  red  colour  in  the  liquid. 

Correction  of  the  Acid.— The  exact  volume  of  the  acid 
which  has  been  used  in  order  to  neutralise  the  25  c.c.  of  normal 
sodium  carbonate  is  noted,  and  from  it  the  volume  of  water  which 
must  be  added  in  order  to  make  the  acid  exactly  normal  is  calcu- 
lated. Thus,  suppose  instead  of  25  c.c.  of  acid,  23'8  c.c.  were  used 
in  neutralising  25  c.c.  of  the  normal  alkali ;  then — 

25  :  23-8  :  :  1000  :  952 

That  is  to  say,  952  c.c.  of  the  acid  contain  as  much  sulphuric 
acid  as  should  be  contained  in  1000  c.c.,  if  it  were  exactly  normal. 
If,  therefore,  952  c.c.  of  this  acid  be  measured  out  into  a  litre-flask, 
and  the  volume  be  then  made  up  to  the  litre  by  the  addition  of 
water,  a  correct  normal  acid  will  be  obtained. 

The  solution  thus  obtained  should  be  once  more  titrated  with 
the  sodium  carbonate. 

When  the  strength  of  the  acid  is  very  nearly,  but  not  exactly, 
normal,  instead  of  attempting  to  bring  it  to  the  precise  normal 
strength,  it  may  be  used  as  it  is,  and  a  correction  introduced  every 
time  by  means  of  a  factor.  For  example,  suppose  25  c.c.  of  normal 
sodium  carbonate  required  24/4  c.c.  of  the  acid  instead  of  25  c.c. 
in  order  to  neutralise  it,  then — 

24-6  :  25 : :  i  :  roi6 

That  is  to  say,  every  cubic  centimetre  of  this  acid  is  equal  to 
roi6  c.c.  of  an  exactly  normal  acid  ;  therefore,  if  the  number  of 

*  Whenever  burettes  or  pipettes  are  employed  for  measuring  standard 
solutions,  they  must  either  be  dry  before  use,  or,  if  moist,  they  must  be  first 
rinsed  out  with  a  small  quantity  of  the  solution  ;  otherwise  that  portion  which 
is  measured  out  for  use  would  be  slightly  diluted  by  the  water  adhering  to  the 
walls  of  the  instrument. 


Acidimetry  and  Alkalimetry.  321 

cubic  centimetres  of  this  acid  used  in  any  titration  be  multiplied 
by  the  factor  roi6,  the  result  will  be  the  number  of  cubic  centi- 
metres which  would  have  been  required  if  the  acid  had  been  strictly 
normal. 

Or  again,  if  the  acid  should  be  a  little  weaker  than  the  exact 
normal,  a  similar  correction  can  be  made.  Thus,  suppose  25  c.c. 
of  normal  sodium  carbonate  required  25-2  c.c.  of  acid  instead  of 
25  c.c.  in  order  to  reach  the  neutral  point,  then — 

25-2:  25  ::  i  10992 

That  is  to  say,  each  cubic  centimetre  of  the  acid  is  in  reality  only 
equivalent  to  0*992  c.c.  of  normal  acid,  therefore  in  this  case  0^992 
is  the  factor  by  which  the  number  of  cubic  centimetres  of  this  acid 
which  might  be  used,  would  have,  to  be  multiplied  in  order  to  con- 
vert them  into  cubic  centimetres  of  normal  acid. 

Titratiou  of  the  Acid,  using  Litmus  as  Indicator. — 
In  the  absence  of  methyl  orange,  the  titration  of  the  acid  by  means 
of  sodium  carbonate  may  be  carried  out  with  litmus  as  the  indi- 
cator. In  this  case,  however,  it  is  necessary  to  boil  the  solution  in 
order  to  expel  the  carbon  dioxide.  The  acid  is  added  gradually, 
until  the  colour  of  the  litmus  changes  from  blue  to  purple-red.  The 
solution  is  then  boiled,  and  as  the  carbon  dioxide  is  expelled  the 
colour  returns  to  the  original  blue  shade.  After  boiling  for  a  few 
minutes,  acid  is  admitted  in  small  quantities,  and  the  liquid  boiled 
up  after  each  addition,  until  at  last  the  addition  of  a  single  drop 
gives  a  permanent  bright-red  colour,  which  does  not  change  on 
boiling. 

As  a  check  to  the  result  thus  obtained,  the  following  method 
may  be  carried  out.  After  noting  the  volume  of  acid  which  has 
been  used  in  order  to  reach  the  point  of  neutrality,  a  measured 
volume  of  acid  in  excess  is  run  into  the  beaker  from  the  burette, 
and  the  mixture  (which  for  this  purpose  is  more  conveniently 
contained  in  a  flask  than  in  a  beaker)  quickly  cooled.  This  excess 
of  acid  is  then  titrated  back  with  normal  caustic  soda,  which 
(since  no  carbon  dioxide  is  evolved)  can  be  carried  out  in  the  cold. 
From  the  amount  of  caustic  alkali  used,  the  actual  volume  of  free 
acid  present  is  ascertained.  If  the  former  titration  was  correct, 
this  should  exactly  agree  with  the  measured  excess  volume  which 
was  added ;  and  if  not,  it  gives  the  amount  of  acid  which,  in  the 
first  operation,  had  been  added  beyond  what  was  actually  necessary 
for  neutrality. 

Control  Experiments.—The  titration  of  the  normal   sul- 


322  Volumetric  Methods  of  Analysis. 

phuric  acid  may  be  checked  by  either  one  or  both  of  the  following 
experiments,  which  should  be  made  in  duplicate— 

(1)  A  small  quantity,  say  3  or  4grams,  of  pure  sodium  bicarbon- 
ate is  heated  in  a  weighed  platinum  crucible  until  the  weight  is 
constant,  as  described  above  in  the  preparation  of  normal  sodium  car- 
bonate solution.     The  contents  of  the  crucible  are  then  dissolved  in 
water,  and  washed  into  a  beaker.    This  solution  is  then  titrated  in  the 
cold  with  the  standard  acid,  using  methyl  orange  as  indicator  ;  or, 
if  litmus  is  employed,  the  titration  is  conducted  in  the  boiling  solu- 
tion.    From  the  result,  the  exact  strength  of  the  acid  is  calculated 
as  follows  :  Suppose  the  weight  of  sodium  carbonate  obtained  after 
heating  the    bicarbonate  was    2 '66   grams,  and   that   this   weight 
required  50  c.c.  of  the  acid  for  complete  neutralisation,  then — 

53  :  2-66  :  :  49  :  2-459  grams  =  weight  of  H2SO4  in  the  50  c.c. 

Then  49  :  2*459  :  \  1000  :  50-19  c.c.  =  volume  of  strictly  normal 
acid  which  contains  2-459  grams  of  H2SO4,  and  which  therefore 
would  be  required  to  neutralise  2'66  grams  of  sodium  carbonate. 
Hence  the  acid  used  is  slightly  above  the  normal  strength.  Its 
factor,  therefore,  will  be  I  -0038  ;  thus — 

50 :  50-19  :  :  i  :  1-0038 

(2)  The  strength  of  the  acid  may  also  be  determined  gravimetri- 
cally  by  taking  50  c.c.  of  the  acid,  and,  after  diluting  with  two  or 
three   times   its   volume  of  water,   precipitating    it   with    barium 
chloride,  and  proceeding  as  described  in  Part  I.,  p.  258. 

Normal  Sodium  Hydroxide. 

(40  grams  of  NaH  O  per  litre. ) 

About  45  grams  of  pure  caustic  soda,  prepared  from  the  metal,* 
are  dissolved  in  water — which  has  been  recently  boiled  to  expel 

*  If  pure  sodium  hydroxide  is  not  to  hand,  it  may  be  prepared  by  either  of 
the  two  following  methods,  namely,  (i)  by  the  action  of  lime  upon  sodium 
carbonate,  or  (2)  by  the  action  of  sodium  upon  water. 

(1)  About  70  grams  of  sodium  carbonate  are  placed  in  a  clean  iron  saucepan 
with  700  c.c.  of  distilled  water,  and  the  mixture  heated  to  boiling.     Forty  grams 
of  good  quicklime  are  made  into  a  paste  with  water,  and  gradually  added  to 
the  boiling  solution,  and  the  mixture  boiled  until  the  whole  of  the  carbonic  acid 
in  the  sodium  carbonate  is  precipkated  as  calcium  carbonate.     The  completion 
of  the  reaction  is  ascertained  by  withdrawing  a  small  quantity  of  the  mixture 
and  allowing  it  to  settle  in  a  test-tube.     If  the  sodium  carbonate  has  been 
wholly  converted  into  caustic,  there  will  be  no  effervescence  -when  a  little  of  the 
clear  liquid  is  acidified  with  a  dilute  acid.     The  saucepan  is  carefully  covered, 
and  the  contents  allowed  to  settle  and  become  cold.     The  clear  liquid  is  then 
decanted  off  into  a  stoppered  bottle,  and  a  portion  of  it  titrated  with  standard 
acid,  as  described  above. 

(2)  A  clean  piece  of  sodium  weighing  about  25  grams  is  cut  into  small  pieces, 
which  are  introduced  one  at  a  time  into  recently  boiled  and  cooled  distilled 
water  in  a  platinum  or  nickel  dish.     A  clock-glass  should  be  placed  over  the 
dish  after  the  addition  of  each  piece  of  sodium,  in  order  to  prevent  the  caustic 


Acidimetry  and  Alkalimetry.  323 

carbon  dioxide,  and  again  cooled — and  the  volume  of  the  liquid 
then  made  up  to  one  litre.  Twenty-five  cubic  centimetres  of  this 
solution  are  then  transferred  to  a  small  flask  by  means  of  a  pipette, 
and  titrated  with  the  standard  sulphuric  acid  in  the  cold,  with 
methyl  orange  as  indicator. 

If  the  caustic  soda  be  free  from  carbonate,  the  "  end  reaction  " 
is  very  sharply  defined  when  litmus  is  used  as  the  indicator  in  the 
cold  ;  but  as  the  sodium  hydroxide,  even  if  free  from  carbonates 
to  begin  with,  is  always  liable  to  absorb  atmospheric  carbon 
dioxide,  it  is  necessary,  when  employing  litmus,  to  boil  off  the 
carbon  dioxide  during  the  titration. 

From  the  results  obtained,  the  volume  of  water  which  must  be 
added  to  the  solution  of  caustic  soda  in  order  to  bring  it  to  normal 
strength  is  calculated  as  explained  in  the  case  of  the  sulphuric  acid 
(p.  320).  When  this  has  been  done,  and  the  liquid  well  shaken  up 
to  ensure  perfect  admixture,  a  fresh  titration  should  be  made  ;  and 
if  the  alkali  is  still  not  absolutely  normal,  the  factor  is  found  by 
which  the  volume  used  is  to  be  multiplied  in  order  to  make  the 
necessary  correction. 

The  standard  sodium  hydroxide  should  be  quickly  transferred 
to  a  store-bottle,  the  stopper  of  which  is  greased  with  a  touch  of 
vaseline ;  and  it  should  be  exposed  as  little  as  possible  to  the  air. 

Normal  Hydrochloric  Acid. 

(36*5  grams  ofRClper  litre.} 

The  ordinary  pure  strong  hydrochloric  acid  contains  28  per  cent, 
of  HC1 ;  40  grams,  therefore,  of  the  real  acid  are  contained  in  143 
grams  of  this  solution.  Since  the  specific  gravity  of  this  aqueous 
acid  is  1*14,  143  grams  will  be  125  c.c.  This  volume,  therefore,  is 
measured  out  into  a  litre  flask,  and  the  volume  made  up  to  the 
1000  c.c  with  water  at  I5°'5. 

The  actual  strength  of  the  acid  so  obtained  is  then  ascertained 
by  titration  with  normal  sodium  hydroxide.  Twenty-five  cubic 

soda  from  being  thrown  out  in  the  event  of  explosions  or  deflagrations  taking 
place.  The  process  of  thus  introducing  the  sodium  in  fragments  which  are 
sufficiently  small  to  render  the  operation  free  from  danger,  is  a  somewhat  long 
and  tedious  operation.  It  may  be  considerably  hastened,  however,  and  the 
metal  exposed  much  less  to  the  action  of  atmospheric  carbon  dioxide,  by  placing 
rather  more  than  the  requisite  quantity  of  sodium  in  a  sodium  wire-press,  and 
slowly  squeezing  the  metal  into  wire,  which  is  received  in  water  in  a  nickel 
or  platinum  dish.  The  fine  wire  of  sodium,  as  it  touches  the  water,  fuses  at 
the  end,  and  is  rapidly  dissolved  in  the  liquid,  and  the  rate  at  which  the  wire 
is  driven  out  can  be  exactly  regulated  so  that  it  shall  be  continuously  fed  into 
the  water. 


324  Volumetric  Methods  of  Analysis. 

centimetres  of  the  alkali  are  transferred  to  a  small  flask,  either  litmus 
or  methyl  orange  being  used  as  the  indicator,  and  the  acid  added 
from  a  burette  until  the  point  of  saturation  is  reached.  From  the 
result  obtained,  the  amount  of  dilution  necessary  to  bring  the  acid 
to  normal  strength  is  calculated. 

When  the  acid  has  been  thus  corrected,  it  is  once  more  titrated 
with  the  caustic  alkali. 

As  a  check  upon  the  result,  the  acid  may  be  titrated  with  the 
normal  sodium  carbonate,  using  methyl  orange  as  indicator,  and 
operating  in  the  cold.  (Litmus  cannot  well  be  used  in  this  case,  as 
the  necessary  boiling  involves  loss  of  hydrochloric  acid.) 

The  exact  strength  of  the  acid  may  also  be  controlled  gravi- 
metrically  by  precipitating  25  c.c.  of  the  liquid  with  silver  nitrate. 

Typical  Analyses  by  means  of  Standard  Acids  and 

Alkalies. 

i.  Estimations  of  the  total  alkali  in  samples  of  caustic 
alkali,  or  in  soda-ash  ;  or  determinations  of  the  acid  in  commercial 
acids  or  acid  liquids,  are  conducted  in  practically  the  same  manner 
as  already  described  for  the  titration  of  the  alkalies  and  acids  used 
for  standard '  solutions.  In  estimating  alkalies  containing  car- 
bonates, if  litmus  is  employed  as  indicator  instead  of  methyl  orange, 
the  "  end  reaction  "  is  rendered  more  certain  if  a  slight  excess  of 
the  normal  acid  is  run  in  from  the  burette,  and  the  mixture  boiled 
to  expel  the  carbon  dioxide.  The  excess  of  acid  is  then  carefully 
titrated  back  with  normal  caustic  soda,  which  is  cautiously  added 
from  a  burette  until  one  drop  restores  the  blue  colour  to  the  litmus. 
The  amount  of  acid  which  is  represented  by  the  volume  of  the 
normal  caustic  soda  thus  used,  is  then  deducted  from  the  original 
total  volume  of  acid  employed,  which  gives  the  exact  amount  of 
acid  required  to  neutralise  the  alkali  under  examination. 

2.  Estimation  of  Caustic  and  Carbonated  Alkali  in 
Soda-ash. — About  5  grams  of  the  soda- ash  are  weighed  into  a  flask, 
and  dissolved  in  water,  and  the  solution  made  up  to  250  c.c.  Fifty 
cubic  centimetres  of  this  solution  are  then  transferred  by  means  of 
a  pipette  to  a  small  flask,  and  the  total  alkali  in  it  determined  by 
titration  with  normal  sulphuric  acid,  using  either  litmus  or  methyl 
orange  as  indicator  (see  above).  The  amount  of  alkali  (calculated 
as  Na2O)  contained  in  50  c.c.  when  multiplied  by  5  (as  only  one- 
fifth  of  the  total  was  used)  gives  the  amount  in  the  original  weight 
taken,  and  from  this  the  percentage  may  be  calculated. 

Another  50  c.c.  of  the  solution  are  next  transferred  to  a  loo-c  c. 
flask,  and  a  solution  of  barium  chloride  added  so  long  as  any 


Acidimetry  and  Alkalimetry.  325 

precipitate  of  barium  carbonate  is  produced.*  The  liquid  is  then 
diluted  with  water  up  to  100  c.c.  and  allowed  to  settle  in  the 
stoppered  flask. 

Fifty  cubic  centimetres  of  the  clear  solution  (equivalent  to  one- 
tenth  of  the  original  weight  taken  for  the  analysis)  are  withdrawn 
with  a  pipette,  transferred  to  a  small  beaker  and  titrated  with  deci- 
normal  hydrochloric  acid  in  the  cold,  with  either  litmus  or  methyl 
orange  as  indicator.  (Sulphuric  acid  cannot  be  employed  on  account 
of  the  barium  present.) 

EXAMPLE. — Weight  of  soda-ash  taken  =  5*05  grams. 
First  Solution. — 50  c.c.  =  one-fifth  total  soda-ash. 
50  c.c.  taken  and  excess  of  normal  sulphuric  acid  added 

(i  c.c.  =  0*031  gram  Na2O).   Volume  of  acid  used...     18*50  c.c. 
Titrated  back  with  normal  caustic  soda  (i  c.c.  =  I  c.c. 

normal  acid).     Volume  required       2*37    „ 

Hence  volume  of  acid  employed  in  neutralising  50  c.c. 

soda-ash  16*13    » 


Therefore  weight  of  Na9O  in  } 

original  weilht  of  soda-ash/  =  l6>13  X  °°3i  x  5  -  2*500  grams 
Percentage  of  Na2O,1       2*50  x  100 
or  total  alkali        }  =     "    5.O5         :  49*55 

Second  Solution. — 50  c.c.  =  one-tenth  total  soda-ash. 
50  c.c.  taken,  and  titrated  with  deci-normal  hydrochloric 
acid  (i  c.c.  =  0*0031  gram  Na2O).     Volume  of  acid 

required      6*3  c.c. 

Therefore  weight  of  Na2Ch 
(present  as   NaHO)  in! 

original  weight  of  soda-    =  6'3  x  °'°°3l  x  lo  =  o>1953  gram 
ash  J 

Hence  the  percentage  oH  _  0*1953  x  100  _ 
Na2O  present  as  caustic  /  ~  ^^ —  °     ' 

Then  total  Na2O  in  original  weight  of  soda-ash  =  2*500 
Na2O  present  as  caustic  alkali  =0-1953 

Therefore  Na2O  present  as  carbonated  alkali   =2*3047  grams 

Hence  the  percentage  of  Na2O  present  asi     2*3047  x  100  . 
carbonate  5-o5 

*  As  soon  as  the  whole  of  the  sodium  carbonate  has  been  decomposed  and 
the  carbonic  acid  precipitated  as  barium  carbonate,  the  barium  chloride  begins 
to  interact  with  the  sodium  hydroxide  present,  forming  sodium  chloride  and 
barium  hydroxide,  the  latter  being  equivalent  to  the  sodium  hydroxide  The 
amount  of  barium  chloride  used,  however,  should  be  as  little  in  excess  of  that 
required  for  the  decomposition  of  the  alkaline  carbonate  as  possible,  for  it  is 
found  in  practice  that  when  barium  hydroxide  is  produced  in  any  quantity  the 
results  obtained  are  too  low. 


326  Volumetric  Methods  of  Analysis. 

The  actual  percentage  of  caustic  alkali  and  carbonated  alkali 
may  readily  be  calculated  ;  thus  —  • 

31  :  40  ::  0*1953  =  0-2474  =  NaHO  equivalent  to  0*1953  grams  Na2O 


Similarly,  31  :  53  :  :  2*3047  =  3*938  =  Na2CO3  equivalent  to 

2'3°47  grams  Na2O 

"*^  77.08  = 


Alternative  Methods.  —  (a)  The  soda-ash  is  dissolved  as 
before,  and  the  total  alkali  determined  in  the  manner  above 
described.  Then  a  fresh  portion  (50  c.c.)  of  the  same  solution  is 
transferred  to  a  flask,  and  the  whole  of  the  carbonate  precipitated 
by  adding  barium  chloride.  Phenolphthalein  is  then  added  for  the 
indicator,  and  the  mixture  at  once  titrated  with  deci-normal  hydro- 
chloric acid  until  the  red  colour  is  discharged.  The  presence  of 
the  precipitated  barium  carbonate  does  not  affect  the  indicator, 
and  the  moment  the  point  is  reached  when  the  caustic  alkali  is 
neutralised,  and  the  standard  acid  begins  to  act  upon  the  barium 
carbonate,  the  liberated  carbon  dioxide  discharges  the  colour  of  the 
indicator. 

(j8)  In  this  method  the  total  alkali  is  not  first  determined.  The 
process  depends  upon  the  fact  that  when  sodium  carbonate  is 
acted  upon  by  acids,  the  neutralisation  takes  place  in  two  stages,  so 
that  when  half  the  acid  necessary  for  complete  saturation  has  been 
added,  the  carbonate  has  been  converted  into  the  bi-carbonate, 
which  has  no  power  to  give  a  red  colour  with  phenolphthalein.  If, 
therefore,  to  a  mixture  of  sodium  hydroxide  and  carbonate  coloured 
red  with  this  indicator,  normal  acid  be  added,  the  phenolphthalein 
will  show  the  point  when  the  whole  of  the  caustic  soda  and  one- 
half  of  the  carbonate  have  become  saturated. 

Fifty  cubic  centimetres  of  the  solution  prepared  as  described 
above  is  coloured  with  phenolphthalein,  and  titrated  with  normal 
sulphuric  acid.  As  soon  as  the  colour  is  discharged,  a  drop  of 
methyl  orange  is  added,  and  the  titration  continued  until  the  yellow 
liquid  becomes  pink. 

The  volume  of  acid  required  in  the  second  stage  (i.e.  during  the 
second  half  of  the  saturation  of  the  carbonate  present)  will  be  one- 
half  of  that  which  has  altogether  been  used  up  in  neutralising  the 
alkaline  carbonate.  Therefore,  if  this  volume  of  acid  be  doubled, 
and  the  result  subtracted  from  the  total  acid  used  in  the  whole 


Acidimetry  and  Alkalimetry.  327 

process,  the  difference  will  represent  the  volume  which  was  used 
in  saturating  the  caustic  alkali  present. 

Thus,  in  the  case  of  the  solution  of  soda-ash  used  in  the  first 
example  (containing  5*05  grams  in  250  c.c.),  5°  c.c.  were  taken,  and 
titrated  with  normal  sulphuric  acid  until  the  phenolphthalein  lost 
its  colour. 

Volume  of  acid  required  '          ...       87    c.c. 

Methyl  orange  was  then  added,  and  the  volume 
cf  acid  required  for  complete  neutralisation          7-45    „ 

Total  acid  used     ...     16-15    » 

Therefore  acid  used  in  saturating  \  _ 

the  carbonate  |  -  7  45  x  2 - 

and  acid  used  in  saturating  J=  _  = 

the  caustic  alkali  j 

These,  when  calculated  as  percentages,  as  in  the  former  example, 
give — 

Percentage  carbonate  =  78' I 
Percentage  caustic  =    4/9 

3.  Estimation  of  Ammonia. 

(a)  Direct  Method. — A  solution  of  the  ammonium  salt  in  water 
is  boiled  with  a  solution  of  caustic  soda,  and  the  ammonia  thus 
expelled  is  absorbed  in  a  measured  volume  of  normal  sulphuric 
acid.  The  excess  of  acid  is  then  ascertained  by  titration  with 
normal  sodium  hydroxide  ;  and  by  deducting  this  excess  from  the 
total  volume  of  acid  employed,  the  volume  of  acid  which  has  been 
saturated  by  the  ammonia  is  found,  and  from  this  the  weight  of 
ammonia  is  calculated. 

From  i  to  2  grams  of  ammonium  chloride  are  weighed  out  into 
the  flask  F  (Fig.  53,  p.  250),  and  dissolved  by  the  addition  of  a 
small  quantity  (40  to  50  c.c.)  of  water.  One  hundred  cubic  centi- 
metres of  normal  sulphuric  acid  are  run  into  the  absorption  flask 
through  the  scrubber-tube,  and  the  process  of  distilling  and  absorb- 
ing the  ammonia  is  carried  out  exactly  as  described  on  p.  251. 

After  the  whole  of  the  ammonia  has  been  expelled,  the  apparatus 
is  disconnected,  and  the  end  of  the  leading  tube  and  the  contents 
of  the  scrubber-tube  are  thoroughly  rinsed  into  the  flask,  so  as  to 
wash  out  of  the  latter  every  trace  of  the  acid  with  which  the  broken 
glass  was  moistened.  A  few  drops  of  methyl  orange  are  then 
added,  and  the  liquid  titrated  with  normal  caustic  soda. 

The  quantity  of  acid  which  is  equivalent  to  the  volume  of  the 
normal  soda  required  for  neutralisation,  is  the  excess  which  has 


328  Volumetric  Methods  of  Analysis. 

remained  over  after  the  whole  of  the  ammonia  has  been  saturated. 
On  deducting  this  from  the  total  volume  of  acid  used  (namely,  100 
c.c.),  the  actual  amount  of  acid  neutralised  by  the  ammonia  is 
found,  and  by  calculation  the  weight  of  ammonia,  NH3  (or  of 
ammonium,  N  H4),  is  ascertained. 

(b)  Indirect  Method. — In  this  method  the  weighed  quantity  of 
ammonium  salt  is  boiled  with  a  measured  volume  of  normal  sodium 
hydroxide,  until  the  ammonia  is  all  expelled.  A  certain  quantity 
of  the  soda  is  thus  decomposed,  being  converted  into  the  sodium 
salt  of  the  acid  which  was  in  combination  with  the  ammonia.  For 
example,  supposing  the  ammonium  salt  to  have  been  the  sulphate — 

(NH4),SO4  +  2NaHO  =  Na2SO4  +  2H2O  +  2NH3 

If,  when  the  whole  of  the  ammonia  has  been  expelled,  the  remain- 
ing undecomposed  caustic  soda  be  determined  by  means  of  normal 
acid,  the  amount  of  soda  which  has  been  used  up  in  expelling  the 
ammonia  will  be  found  ;  and  from  this  the  quantity  of  ammonia  so 
displaced  can  be  calculated. 

About  2  grams  of  pure  ammonium  sulphate  are  weighed  out 
into  a  small  flask,  and  30  c.c.  of  normal  caustic  soda  added.  The 
mixture  is  then  boiled  until  the  escaping  steam  produces  no  brown 
stain  upon  a  strip  of  turmeric  paper.  Three  or  four  drops  of 
litmus  solution  are  then  added,  and  the  liquid  titrated  with  normal 
sulphuric  acid.  The  volume  of  normal  soda  equivalent  to  the 
number  of  cubic  centimetres  of  acid  used  to  produce  neutrality, 
when  deducted  from  the  total  volume  of  soda  employed  (namely, 
30  c.c.)  gives  the  number  of  cubic  centimetres  of  soda  which  were 
decomposed  in  the  process  of  expelling  the  ammonia. 

From  the  above  equation,  it  is  seen  that  40  grams  of  caustic 
soda  are  used  up  in  the  expulsion  of  17  grams  of  ammonia  ;  there- 
fore each  cubic  centimetre  of  normal  soda  (0*040  gram)  is  equivalent 
to  o'oiy  gram  of  ammonia,  NH3. 

4.  Estimation  of  the  "Hardness"  of  Water*  (Hehner's 
method). 

(a)  Temporary  hardness  (due  to  the  presence  of  calcium  or 
magnesium  carbonate)  is  determined  by  means  of  standard  sul- 
phuric acid  of  one-fiftieth  normal  strength.  This  is  prepared  by 
diluting  20  c.c.  of  the  normal  acid  up  to  1000  c.c.  with  water.  One 
cubic  centimetre  of  this  acid  is  equivalent  to  o'ooi  gram  CaCO3. 

One  hundred  cubic  centimetres  of  the  water  to  be  examined  are 
transferred  to  a  beaker,  and  3  drops  of  methyl  orange  added. 

*  For  Clark's  soap-test  for  hardness  in  water,  see  Precipitation  methods, 
P-  37i. 


Acidimelry  and  Alkalimetry.  329 

The  —  sulphuric  acid  is  then  delivered  from  a  burette  until  the 

50 
solution  becomes  red. 

Each  cubic  centimetre  of  acid  represents  i  milligram  of  CaCO3 
in  100  c.c.  of  water  ;  therefore  it  also  stands  for  parts  per  100,000. 
E.g.  suppose  13  c  c.  of  acid  were  required  to  exactly  neutralise  the 
carbonate  ;  then  100  c.c.  of  the  water  contain  0*013  gram  of  CaCO3 ; 
or  its  equivalent  of  MgCO2,  and  100,000  parts  contain  13  parts  of 
CaCO3  (or  the  equivalent  of  MgCO3). 

One  part  of  CaCO3  per  100,000  of  water  is  termed  I  degree  of 
hardness  ;  hence  the  water  in  question  has  a  temporary  hardness 
of  1 3  degrees. 

(b}  Permanent  hardness  (due  mainly  to  the  presence  of  calcium 
or  magnesium  sulphate)  is  estimated  by  means  of  standard  sodium 
carbonate  of  one-fiftieth  the  normal  strength  (prepared  by  diluting 
20  c.c.  of  the  normal  solution  to  i  litre  with  water).  One  cubic 
centimetre  of  this  solution  will  precipitate  o'ooi  gram  CaCO3. 
Excess  of  this  is  added  to  the  water,  which  is  then  evaporated  to 
dryness.  The  residue  is  treated  with  water,  and  after  removing 
the  precipitated  carbonates  by  filtration,  the  excess  of  sodium 

N 
carbonate  is  determined  by  means  of  —  sulphuric  acid. 

One  hundred  cubic  centimetres  of  the  water  are  transferred  to 

N 
a  platinum  dish,  and  a  measured  volume  of  the  —  sodium  carbonate 

— considerably  in  excess  of  that  required  for  complete  precipitation 
— is  added.  The  solution  is  then  evaporated  to  dryness  on  a  steam 
bath.  The  residue  is  taken  up  with  water  and  filtered,  the  dish 
and  filter  being  thoroughly  washed.  The  sodium  carbonate  in  the 

N 
liquid  is  then  determined  by  titration  with  the   —  sulphuric  acid. 

The  excess  of  sodium  carbonate  thus  found,  deducted  from  the 
total  volume  added,  gives  the  number  of  cubic  centimetres  used  in 
removing  the  permanent  hardness.  Each  cubic  centimetre  so 
employed  represents  o'ooi  gram  of  calcium  carbonate  (formed  from 
the  sulphate  *)  precipitated  from  100  c.c.  of  water  ;  it  therefore 
also  represents  parts  per  100,000. 

*  From  the  following  equations,  it  will  be  seen  that,  although  sodium 
carbonate  precipitates  calcium  carbonate  from  the  soluble  bicarbonate  of  lime 
(temporary  hardness)  as  well  as  from  the  sulphate  (permanent  hardness) — that, 
in  other  words,  it  removes  both  temporary  and  permanent  hardness — it  is  only 
in  decomposing  the  sulphate  that  the  alkali  is  destroyed.  In  the  other  case 
the  solution  contains  the  equivalent  of  sodium  bicarbonate — 

H2CO3,CaCO3  4-  Na2CO3  =  CaCO3  +  2HNaCO3 
CaS04  +  Na2C03  =  CaCO3  +  Na2SO4 


330  Volumetric  Methods  of  Analysis, 

N 
For  example,  100  c.c.  of  the  water  were  taken,  and  50  c.c.  of  - 

sodium  carbonate  added.     After  evaporating  to  dryness  and  ex- 

N 
tracting  with  water,  the  solution  required  41  c.c.  of  —  sulphuric 

acid  to  neutralise  it.  50  —  41  =  9,  therefore  9  c.c.  of  the  sodium 
carbonate  were  used  up  in  removing  the  permanent  hardness, 
hence  the  water  contained  9  degrees  of  permanent  hardness.  The 
temporary  hardness  added  to  the  permanent  hardness  gives  the 
total  hardness  of  the  water. 


SECTION    III. 

VOLUMETRIC  METHODS  BASED  UPON  OXIDATION  AND 
REDUCTION. 

IN  the  various  processes  which  are  included  in  this  section,  there 
is  always  an  oxidising  agent  (in  other  words,  a  reducible  substance) 
and  an  oxidisable  body,  i.e.  a  reducing  agent. 

Sometimes  the  oxidising  material  is  the  standard  reagent, 
which,  by  its  own  reduction,  furnishes  the  measure  of  the  amount 
of  the  oxidisable  substance  which  is  present.  Thus,  the  reduction 
of  a  known  volume  of  standard  potassium  permanganate,  or  potas- 
sium dichromate,  will  afford  a  measure  of  the  quantity  of  iron 
which  has  been  oxidised  from  the  ferrous  to  the  ferric  condition  ; 
the  ferrous  salt  in  this  case  being  the  oxidisable  substance,  or  the 
reducing  agent. 

In  other  cases  the  standard  reagent  is  the  reducing  agent, 
which  undergoes  oxidation  at  the  expense  of  some  oxidising  (i.e. 
reducible)  compound  present,  and  thereby  affords  the  means  of 
estimating  the  amount  of  such  a  substance.  Thus,  the  oxidation 
of  a  known  volume  of  standard  sodium  arsenite  into  arsenate,  may 
become  the  measure  of  the  amount  of  chlorine  present  in  bleach- 
ing powder ;  or  of  the  quantity  of  free  iodine  in  a  solution.  The 
chlorine  or  the  iodine  in  this  case  is  the  oxidising  agent,  not  by 
directly  giving  oxygen,  as  in  the  case  of  potassium  permanganate, 
but  by  indirect  oxidation  through  the  intervention  of  water  ;  the 
hydrogen  of  which  is  taken  by  the  halogen,  while  an  equivalent 
of  oxygen  is  liberated. 

The  most  important  of  the  oxidising  standard  solutions  are 
potassium  permanganate,  potassium  dichromate,  and  iodine  ;  while 
the  reducing  compounds  in  commonest  use  are  ferrous  salts,  oxalic 
acid,  sodium  thiosulphate,  sodium  arsenite. 


332  Volumetric  Analysis. 

Deci-normal  Potassium  Permanganate. 

(3-16  grams  of  KMnO4  per  litre*} 

3*16  grams  of  the  purest  recrystallised  and  dry  salt  f  are  weighed 
out  and  dissolved  in  water  in  a  litre  flask,  and  the  solution  made 
up  to  1000  c.c.  with  water  at  15°. 

Solutions  of  potassium  permanganate  slowly  undergo  decom- 
position by  light ;  if,  therefore,  the  solution  is  to  be  kept  for  any 
length  of  time,  it  is  preferably  transferred,  either  to  a  deep-blue 
bottle,  or  to  one  covered  with  black  paper,  and  kept  in  a  cupboard. 
The  solution  must  not  be  allowed  to  come  in  contact  with  any 
organic  matter  ;  on  this  account,  burettes  with  glass  taps  must 
always  be  employed  for  this  reagent. 

Titration  of  Deci-normal  Permanganate.— (a)  By  means 
of  Ferrous  Sulphate — 

2KMn04  +  ioFeSO4  +  8H2SO4  =  K2SO4  +  2MnSO4  +  8H2O 

+  5Fea(S04)3 

or,  expressing  only  the  relation  between  the  available  oxygen  and 
the  ferrous  oxide — 

50  +  icFeO  =  5Fe2O3 

The  deci-normal  solution  being  one-tenth  of  the  strength  of  the 
normal  solution — which,  in  its  turn,  contains  (in  the  case  of  per- 
manganate) only  one-tenth  of  the  formula-weight  of  the  salt  per 
litre — will  contain  o'oooS  gram  of  available  oxygen  per  cubic  centi- 
metre ;  or  each  cubic  centimetre  will  be  equivalent  to  0-0056  gram 
of  iron  in  the  ferrous  state  ;  i.e.  it  will  be  capable  of  oxidising  this 
weight  of  iron  from  tint  ferrous  to  the  ferric  condition. 

For  the  purposes  of  this  titration,  the  ferrous  sulphate  is  pre- 
pared by  dissolving  o-5  gram  of  the  purest  soft  iron  wire  in 
sulphuric  acid,  with  exclusion  of  air,  the  wire  being  clean  and  free 
from  rust.t 

About  80  c.c.  of  dilute  sulphuric  acid  (i  part  acid  to  5  parts 
water)  are  placed  in  a  25o-c.c.  flask  fitted  with  a  rubber  cork  and 
bent  glass  tube.  The  air  in  the  flask  is  then  expelled  by  removing 
the  cork  and  introducing  two  or  three  crystals  of  pure  sodium 
carbonate,  the  flask  being  in  a  vertical  position.  As  soon  as  the 
carbonate  has  dissolved,  the  weighed  quantity  of  iron  is  dropped 

*  See  p.  314. 

f  If  the  ordinary  "  pure"  commercial  salt  be  used,  about  3'2  to  3^5  grams 
should  be  taken. 

\  The  fine  iron  binding  wire  used  for  flowers  is  the  best  for  the  purpose  ;  it 
contains  99*6  per  cent,  of  iron. 


Oxidation  and  Reduction  Methods. 


333 


in.  The  cork  is  then  inserted,  and  the  flask  supported  in  the 
manner  shown  in  Fig.  72,  with  the  tube  dipping  into  a  solution  of 
sodium  carbonate  in  a  small  beaker.  The  flask  being  in  this  in- 
clined position,  the  fine  spray  thrown  up  during  the  solution  of  the 
iron  strikes  against  the  sides  of  the  flask  and  falls  back  into  the 
liquid.  The  flask  is  gently 
heated  by  means  of  a  small 
flame  until  the  iron  is  wholly 
dissolved,  and  only  a  few  minute 
particles  of  carbon  remain.  The 
lamp  is  then  withdrawn,  and  the 
flask  allowed  to  cool.  As  it  does 
so,  the  solution  in  the  beaker  is 
gradually  drawn  up  the  tube, 
but  the  first  drops  which  enter 
the  flask  at  once  cause  an  effer- 
vescence of  carbon  dioxide 
which  drives  the  liquid  down 
again,  and  at  the  same  time  fills 
the  flask  with  carbon  dioxide. 
When  it  has  partially  cooled  in 
this  way,  the  cork  is  removed, 
and  air-free  distilled  water  (pre- 
pared by  boiling  the  water,  and 
again  quickly  cooling  it)  is 
added  until  the  solution  is  within 
about  20  or  30  c.c.  of  the  gra- 
duation mark.  The  flask  is  then 
closed  with  a  rubber  stopper,  and  the  contents  made  quite  cold  by 
holding  the  vessel  in  a  stream  of  cold  water.  The  solution  is  then 
made  up  to  250  c.c.  by  the  further  addition  of  cold  air-free  water. 

Fifty  cubic  centimetres  of  this  solution  are  transferred  by  means 
of  a  pipette  to  a  small  flask,  and  diluted  by  the  addition  of  about 
half  the  volume  of  air-free  distilled  water.  The  flask  is  placed 
upon  a  white  tile,  and  the  deci-normal  permanganate  solution  added 
from  a  burette  until  the  colour  of  the  reagent  ceases  to  be  destroyed, 
and  a  faint  pink  tint  is  imparted  to  the  solution. 

Four  separate  experiments  should  be  made,  taking  50  c.c.  of  the 
iron  solution  each  time,  in  order  to  gain  practice  in  judging  when 
the  first  appearance  of  the  permanent  pink  colour  takes  place. 
After  the  experience  thus  gained,  in  subsequent  duplicate  titrations 
the  volume  of  the  reagent  used  should  agree  to  o-i  of  a  cubic 


FIG.  72. 


334  Volumetric  Analysis. 

centimetre.  From  the  results  obtained,  the  exact  strength  of  the 
permanganate  is  calculated  ;  thus  — 

o"5  gram  of  iron  wire  was  dissolved  in  250  c.c.  of  liquid. 
Fifty  cubic  centimetres  of  the  solution  therefore  contain  OT  gram 
of  iron.  But  since  the  iron  wire  contained  99-60  per  cent,  of  iron, 
the  actual  weight  of  iron  present  is  found  by  multiplying  the 
weight  of  the  wire  by  0-996.  Hence  50  c.c.  of  the  solution  contain 
not  0*1  gram,  but  0-0996  grams  of  iron.  Three  titrations  were 
made,  which  required  respectively  17*64  c.c.,  17-74  c-c.,  and  17-70 
c.c.  of  the  permanganate,  or  17*69  as  a  mean  figure. 

Then  17*69  c.c. :  i  c.c.  : :  0-0996  gram  :  0*00563  gram 

Therefore  the  solution  is  very  slightly  stronger  than  the  exact 
deci-normal,  since  i  c.c.  should  be  equivalent  to  0-0056  gram  of 
iron. 

(f)  Titration  of  Permanganate  with  Ferrous  Ammonium  Std- 
phate,  FeSO4,(NH4)2SO4,6H2O.— The  formula-weight  of  this  salt 
(392)  is  exactly  seven  times  the  weight  of  the  iron  it  contains  (56)  ; 
hence  in  3*5  grams  of  the  compound  there  will  be  0-5  gram  of  iron. 

Exactly  3-5  grams  of  the  pure  salt  are  weighed  out  and  dissolved 
in  air-free  water,  and  the  solution  made  up  to  250  c.c.  Fifty  cubic 
centimetres  of  this  solution,  therefore,  will  contain  o'i  gram  of  iron. 
For  the  deration,  50  c.c.  of  the  solution  are  transferred  to  a  small 
flask,  and  10  or  15  c.c.  of  dilute  sulphuric  acid  added.  The  deci- 
normal  solution  is  then  run  in,  as  in  the  former  case,  until  the  pink 
colour  persists. 

(c]  Titration  of  Permanganate  with  Oxalic  Acid,  H2C2O4,2H2O. 

2KMnO4  +  5H2C204  +  3H2SO4  =  K2SO4  +  2MnSO4  +  ioCO2 

+  8H20 
or  50  +  5H2C2O4  =  5H2O  +  ioCO2 

Hence  8  parts  of  oxygen  (i  equivalent)  is  capable  of  oxidising  45 
parts  of  oxalic  acid  (calculated  as  anhydrous),  or  63  parts  of  the 
crystallised  compound.  Therefore  i  c.c.  of  the  deci-normal  per- 
manganate, which  is  equivalent  to  0*0056  gram  of  iron,  will  be 
equivalent  also  to  0-0063  gram  of  crystallised  oxalic  acid. 

6*3  grams  of  recrystallised  and  dry  *  oxalic  acid  are  weighed 
out,  dissolved  in  water,  and  the  solution  diluted  to  i  litre.  The 
solution  so  obtained  is  deci-normal  oxalic  acid  ;  but,  in  order  to  make 
sure  that  the  acid  is  pure,  and  of  the  exact  state  of  hydration  repre- 
sented by  the  formula  C2H2O4,2H2O,  the  precise  strength  of  the 

*  The  term  dry  signifies  freedom  from  adherent  moisture,  and  does  not 
refer  to  the  anhydrous  compound. 


Oxidation  and  Reduction  Methods.  335 

solution  should  be  determined  by  titration  with  adeci-normal  caustic 
alkali,  using  phenol phthalein  as  indicator. 

Twenty-five  cubic  centimetres  of  the  deci-normal  oxalic  acid  are 
transferred  to  a  small  flask,  and  about  half  the  volume  of  dilute 
sulphuric  acid  added.*  The  mixture  is  then  gently  heated  to  a 
temperature  from  50°  to  60°,  and  the  permanganate  solution 
slowly  run  in  from  a  burette.  In  this  reaction  the  discharge  of 
the  colour  of  the  permanganate  is  not  so  instantaneous  as  in  the 
case  of  the  ferrous  compounds,  there  being  an  intermediate  forma- 
tion of  a  brown  colour,  which  then  fades  away,  leaving  the  liquid 
colourless.  Hence  it  is  necessary  to  add  the  permanganate 
cautiously,  especially  as  the  end  of  the  reaction  approaches,  in 
order  that  the  first  signs  of  the  permanent  pink  colour  may  be 
detected.  Two  or  three  titrations  should  be  made  (using  various 
quantities  of  the  deci-normal  oxalic  acid)  for  practice  in  observing 
the  end  reaction. 

If  the  oxalic  acid  be  strictly  deci-normal  (as  determined  by 
means  of  deci-normal  alkali), -and  the  permanganate  be  likewise 
exact,  25  c.c.  of  the  acid  should  require  25  c.c.  of  permanganate. t 

Typical  Analyses  by  means  of  Potassium  Perman- 
ganate. 

(i)  Estimation  of  Iron  in  Iron  Ores.J — If  the  iron  is 
present  in  the  ferrous  state,  the  substance  is  dissolved  with  careful 
exclusion  of  air,  and  the  solution  titrated  as  already  described. 
If  the  iron  is  in  the  ferric  condition,  it  is  reduced  (after  solution) 
by  the  method  given  below.  When  the  metal  is  present  partly  as 
ferrous  and  partly  ^ferric  iron,  the  ferrous  iron  is  first  estimated 
in  one  portion,  and  a  second  portion  is  then  subjected  to  the 
reducing  agent,  and  the  total  amount  of  iron  determined.  The 
difference  between  the  two  results  represents  the  iron  which  existed 
in  the  ferric  state. 

(a]  Iron  in  the  Ferrous  State. — About  0*5  gram  of  spathic  iron 
ore,  in  \^\o.  finest  possible  powder^  is  weighed  out  and  dissolved  in 

*  The  presence  of  acid  (preferably  sulphuric  acid)  is  necessary  in  all  titra- 
tions with  potassium  permanganate,  in  order  to  keep  the  manganous  oxide  in 
solution.  Nitric  and  hydrochloric  acids  are  prejudicial ;  with  the  latter  acid 
(unless  present  only  in  small  quantities  and  the  liquid  is  cold),  chlorine  is 
liberated  by  the  oxidising  action  of  the  permanganate. 

f  Dilute  solutions  of  oxalic  acid,  such  as  the  deci-normal,  are  liable  to 
undergo  decomposition  if  preserved  ;  hence,  if  the  solution  has  been  prepared 
any  length  of  time,  it  should  be  retitrated  with  alkali  before  being  used. 

J  Such  as  magnetic  oxide,  red  or  brown  hcematite,  clay  ironstone,  spathose, 
black  band. 


336  Volumetric  Analysis. 

20  c.c.  of  a  mixture  of  equal  volumes  of  strong  hydrochloric  acid 
and  water,  with  exclusion  of  air,  in  the  apparatus  described  on  p. 
333.*  The  acid  is  gently  boiled  by  means  of  a  small  flame  until 
the  iron  is  entirely  extracted.  The  solution  is  then  largely  diluted 
by  adding  100  c.c.  of  cold  air-free  water,  and  at  once  titrated  by 
means  of  deci-normal  permanganate.f 

(£)  Total  Iron. — A  second  portion  of  ore,  about  0*5  gram,  is 
weighed  out  and  dissolved  in  a  similar  quantity  of  acid.  The 
operation  may  be  conducted  in  a  flask  supported  as  shown  in  Fig. 
72,  but  the  cork  with  the  leading  tube  may  be  removed. 

When  the  solution  of  the  ore  is  complete,  the  liquid  is  con- 
siderably diluted  with  water,  and  about  15  c.c.  of  strong  sulphuric 
acid  added.  A  few  fragments  of  pure  zinc  (i.e.  free  from  ironj) 
are  then  introduced  into  the  flask,  and  the  cork  with  its  leading 
tube  is  inserted.  Hydrogen  is  evolved  by  the  solution  of  the  zinc 
in  the  acid  liquid,  and  the  iron  existing  in  the  ferric  state  is  thereby 
reduced  to  the  ferrous  condition.  The  action  is  allowed  to  continue 
until  the  zinc  is  entirely  dissolved,  the  process  being  aided  towards 
the  end  by  the  application  of  heat.  In  order  to  test  whether  the 
reduction  of  the  iron  is  complete,  a  drop  of  the  liquid  is  with- 
drawn upon  the  end  of  a  fine  glass  rod,  and  brought  into  contact 
with  a  drop  of  a  solution  of  ammonium  thiocyanate  upon  a  white 
tile. 

Any  remaining  ferric  salt  will  be  revealed  by  the  formation  of 
the  red  colour ;  in  which  case  the  process  must  be  continued  by 
the  addition  of  more  zinc,  and,  if  necessary,  of  more  acid  also. 
When  the  iron  is  entirely  in  the  ferrous  state,  the  solution  is  quickly 
cooled,  and  then  titrated  with  the  permanganate. 

By  deducting  the  amount  of  iron  present  as  ferrous  iron  from 
the  total  iron  in  the  ore,  the  amount  of  iron  which  is  in  the  ferric 
state  is  ascertained. 

*  In  the  case  of  the  carbonate  of  iron,  the  carbon  dioxide  evolved  will  serve 
to  displace  the  air  without  the  addition  of  sodium  carbonate. 

f  It  is  very  necessary  that  the  amount  of  hydrochloric  acid  present  should 
be  quite  small,  and  the  solution  cold  and  dilute,  otherwise  the  analysis  will  be 
yitiated  by  the  action  of  the  acid  upon  the  permanganate,  whereby  chlorine  is 
liberated;  thus— 

KMnO4  +  8HC1  =  KC1  +  MnCl2  +  4H2O  +  sCl 

|  If  pure  zinc  is  not  available,  the  zinc  that  is  employed  must  be  weighed 
before  being  introduced,  and  the  amount  of  iron  present  must  be  found  by  dis- 
solving a  similar  quantity  of  the  zinc  in  acid,  and  titrating  with  permanganate. 
Or  magnesium  may  be  used  instead  of  zinc.  Other  methods  by  which  iron 
may  be  reduced,  but  which  are  more  applicable  when  the  titration  is  to  be  made 
with  potassium  dichromate,  are  described  on  pp.  343,  344. 


Analyses  by  means  of  Potassium  Permanganate.     337 

EXAMPLE. — 0*5  gram  of  spathic  iron  ore. 
(a)  Ferrous  Iron. 

Permanganate  required 34*4  c.c. 

(i  c.c.  =  0*0056  Fe,  =  0*0072  FeO,  =  0*0080  Fe2O3) 

Then  34*4  x  0*0072  =  0*24768  =  weight  of  FeO 

and  °'24768  X  I0°  =  49*53  =  per  cent,  of  FeO  in  the  ore 

(£)  Ferric  Iron. — 0*5  gram  of  ore. 

Permanganate  required        36*6  c.c. 

Deducting  the  volume  used  in  the  first  titration 34*4   „ 

Gives  the  volume  required  for  the  iron  present  in  the 

ferric  state  2*2   „ 

But  the  zinc  used  contained  0*08  per  cent,  of  iron.  Five  grams  were 

employed  ;   therefore  -          -  =  0*004  gram    of  iron  which   was 

introduced. 

The  volume  of  permanganate  used  in  oxidising  this 

quantity  of  iron  is  0-0056  :  0*0040  : :  i  :  071  c.c.          2*2 

><s  071 

Deducting  this,  gives  the  true  volume  of  permanganate 
required  for  the  iron  originally  present  as  ferric  iron         1*49  C.C. 

Then  1*49  x  0*0080  =  0*01192  =  weight  of  Fe2O3 

.  0*01192  x  100 
and  —     — =  2*38  =  per  cent,  of  Fe2O3  in  the  ore 

(2)  Estimation  of  Available   Oxygen  in  Manganese 

Ores. — The  available  oxygen  in  a  manganese  ore  is  that  oxygen 
which  can  be  made  use  of  for  oxidising  purposes  when  the  ore  is 
decomposed  by  an  acid.  Thus,  in  the  sesquioxide  one-third,  and 
in  the  dioxide  one-half,  of  the  total  oxygen  is  available  for  this  pur- 
pose. When  acted  upon  by  sulphuric  acid,  this  oxygen  is  evolved 
as  such  j  thus — 

Mn2O3  +  2H2SO4  =  2MnSO4  +  2H20  +  O 
MnO2  +  H2SO4  =  MnSO4  +  H2O  +  O 

while,  when  decomposed  by  hydrochloric  acid,  chlorine  is  evolved, 
in  each  case  in  quantity  equivalent  to  the  oxygen — 

Mn2O3  +  6HC1  =  2MnCl2  +  3H2O  +  C12 
MnO2  +  4HC1  =  MnCl2  +  2H2O  +  C12 

Methods  for  valuing  manganese  ores  are  based  upon  both  of 
these  reactions.     In  the  first  case  the  evolved  oxygen  is  used  in 


338  Volumetric  Analysis. 

oxidising  either  a  ferrous  salt  or  oxalic  acid  ;  while  in  the  second, 
the  chlorine  is  made  to  liberate  its  equivalent  of  iodine,  which  is 
then  determined  by  means  of  sodium  thiosulphate. 

In  this  place,  examples  based  upon  the  first  reaction  are  given  ; 
the  second  method  is  described  on  p.  356. 

(a)  By  the  Oxidation  of  Ferrous  Stdphate. 

MnO2  +  2H2SO4  +  2FeSO,  =.  Fe2(SO4)3  +  MnS04  +  2H2O 

That  is  to  say,  i  atom  or  16  parts  of  oxygen  (which  is  the 
representative  of  I  equivalent  (or  87  parts)  of  manganese  dioxide), 
will  oxidise  2  atoms  or  1  12  parts  of  ferrous  iron. 

If,  therefore,  manganese  dioxide  be  dissolved  in  sulphuric  acid, 
in  the  presence  of  a  known  quantity  of  ferrous  salt  (which  must  be 
in  excess),  every  8  parts  of  oxygen,  or  43-5  parts  of  MnO2  will  oxidise 
56  parts  of  Fe  ;  and  if  the  amount  of  ferrous  salt  which  remains 
unoxidised  be  then  determined  by  means  of  deci-normal  perman- 
ganate, the  quantity  of  iron  oxidised,  and  therefore  the  amount 
of  manganese  dioxide,  can  be  calculated. 

About  i  gram  of  soft  iron  wire  is  weighed  out  and  dissolved  in 
about  40  c.c.  of  dilute  sulphuric  acid  (i  part  acid  to  4  parts  water) 
in  the  apparatus  shown  on  p.  333. 

As  soon  as  it  is  completely  dissolved,  075  to  i  gram  of  finely 
powdered  pyrolusite,  previously  dried  at  100°,  is  introduced  into 
the  flask,  and  the  cork  with  its  tube  replaced.  The  mixture  is 
gently  heated  until  the  ore  has  entirely  dissolved,  or  until  all  visible 
black  particles  have  disappeared.  The  solution  is  cooled  in  the 
manner  described  on  p.  333,  diluted  with  air-free  water,  and  the 
excess  of  ferrous  salt  titrated  with  permanganate. 

The  excess  of  unoxidised  iron  thus  found,  deducted  from  the 
weight  of  iron  taken,  gives  the  quantity  of  iron  which  has  been 
oxidised  by  the  manganese  dioxide  ;  and  from  this,  by  the  equation 
given  above,  the  percentage  of  available  oxygen  —  either  as  oxygen, 
or  in  terms  of  manganese  dioxide  —  can  be  calculated.  Thus  — 

Volume  of  permanganate  \ 

(i  c.c.  =  0-0056  Fe)  required/  ~ 

22*4  x  0-0056  =  o'i25  gram  =  excess  of  Fe 


("'6  ^  }  =  °'996  gram  =  total  iron  taken 
therefore  0-996  -  0-125  =  0-87  1  gram  =  iron  oxidised  by 
075  gram  of  the  ore 

Fe          Fe  MnO2 

Then  56  :  0*871  :  :  43-5  :  0-6765  gram  MnO2 


Analyses  by  means  of  Potassium  Permanganate.     339 


and  I00=  90  o  -percentage  of  MnO2  in  the 

°75  ore 

(b]  By  the  Oxidation  of  Oxalic  Acid— 

MnO2  -f  H2SO4  +  C2H2O4  =  MnSO,  +  2CO2  +  2H2O 

Hence  every  8  parts  of  oxygen,  or  43-5  parts  of  MnO2,  will  oxidise 
63  parts  of  crystallised  oxalic  acid.  About  0-25  gram  of  the 
manganese  ore  is  weighed  into  a  flask,  and  10  c.c.  of  normal  oxalic 
acid  added.  Twenty-five  cubic  centimetres  of  dilute  sulphuric  acid 
(i  part  acid  to  4  parts  water)  are  added,  and  the  mixture  gently 
warmed  until  the  decomposition  is  complete  and  no  black  particles 
are  visible.  The  solution  is  then  cooled,  and  made  up  to  100  c.c. 
with  cold  water.  The  excess  of  oxalic  acid  remaining  is  titrated 
with  deci-normal  permanganate,  taking  25  c.c.  of  the  solution. 
The  amount  of  oxalic  acid  thus  found,  deducted  from  the  total 
employed,  gives  the  quantity  which  has  been  oxidised  by  the 
manganese  ore,  from  which  the  percentage  of  available  oxygen 
or  of  manganese  dioxide  can  be  calculated.  For  example,  after 
acting  upon  0*25  gram  of  pyrolusite  — 

25  c.c.  of  the  solution  required  of  permanganate  (i  c.c. 

=  0*0063  gram  oxalic  acid)      ............     I2'l  c.C. 

Multiplying  by  4  gives  for  the  total  excess  of  oxalic  acid      48*4  „ 

48*4  x  0-0063  =  weight  of  oxalic  acid  in  excess  =  0-3049  gram 

Weight  of  oxalic  acid  in  10  c.c.  normal  solution-..     0-63 
Deducting  the  excess  ............    o'3049 

Gives  the  weight  of  oxalic  acid  oxidised  by  the  ore    0-325  1 
Then  63  :  0-3251  :  :  43-5  =  0-223  =  weight  of  MnO2 

,  0223  x  loo 
and  --  —  ---  =  89  2  =  percentage  of  MnO2 

(3)  Estimation  of  Tin.  —  When  stannous  chloride  is  mixed 
with  ferric  chloride,  or  when  metallic  tin  is  dissolved  in  ferric 
chloride  mixed  with  hydrochloric  acid,  the  stannous  chloride  is 
oxidised  to  stannic  chloride  at  the  expense  of  the  ferric  chloride. 
A  quantity  of  ferrous  iron,  equivalent  to  the  weight  of  tin,  is  thus 
formed  in  the  solution  according  to  the  equations  — 

Fe,Cl0  +  SnCl2  =  SnCl,  +  2FeCl2 
2Fe2Cl6  +  Sn  =  SnCl4  +  4FeCl2 

Every  56  parts  of  iron  converted  into  the  ferrous  state,  therefore, 
is  equivalent  to  59  or  29*5  parts  of  tin  respectively.  The  amount 
of  ferrous  salt  so  produced  is  titrated  by  means  of  potassium 
permanganate. 


346  Volumetric  Analysis. 

About  O'5  gram  of  granulated  tin  is  weighed  into  a  250-0.0. 
flask,  and  a  few  cubic  centimetres  of  a  moderately  strong  solution 
of , ferric  chloride,  together  with  a  little  hydrochloric  acid,  are  added. 
A  crystal  of  sodium  carbonate  is  dropped  into  the  mixture  in  order 
to  replace  the  air  with  carbon  dioxide,  and  the  flask  loosely  corked 
and  gently  warmed.  When  the  tin  has  wholly  dissolved,  the 
solution  (still  yellow  with  excess  of  ferric  chloride)  is  diluted  up  to 
250  c.c.  with  cold  air-free  water.  Fifty  cubic  centimetres  of  this 
liquid  are  then  withdrawn  by  means  of  a  pipette,  and  titrated  with 
deci-normal  permanganate.* 

(4)  Estimation  of  Calcium. 

(a)  The  calcium  is  precipitated  in  the  form  of  calcium  oxalate 
by  means  of  ammonium  oxalate,  in  the  manner  described  for  the 
gravimetric  estimation,  p.  231.  The  precipitate,  after  being 
thoroughly  washed  to  remove  every  trace  of  ammonium  oxalate, 
is  dissolved  in  a  small  quantity  of  warm  hydrochloric  acid,  and  the 
solution  diluted  with  water.  Sulphuric  acid  is  then  added,  and  the 
mixture  heated  to  between  60°  and  70°,  and  the  free  oxalic  acid 
titrated  by  means  of  permanganate.  Sixty-three  parts  of  oxalic 

N 
acid  found,  are  equivalent  to  20  parts  of  calcium  ;  or   I   c.c.   — 

permanganate,  which  is  equal  to  0*0063  gram  oxalic  acid,  represents 
O'oo2  gram  of  calcium. 

About  0-5  gram  of  marble  or  limestone  is  weighed  out,  and  dis- 
solved and  precipitated  by  means  of  ammonium  oxalate  in  the 
manner  described  on  p.  231.  The  perfectly  washed  precipitate  is 
then  washed  into  a  flask  by  thrusting  a  glass  rod  through  the  apex 
of  the  filter,  and  that  which  remains  adhering  to  the  paper  is  dis- 
solved off  by  treating  it  with  a  little  warm  hydrochloric  acid.  The 
filter  is  thoroughly  rinsed  with  water,  and  if  the  whole  of  the 
precipitate  in  the  flask  is  not  dissolved,  a  few  more  drops  of  acid 
are  added,  avoiding  unnecessary  excess.  A  few  cubic  centimetres 
of  strong  sulphuric  acid  are  now  added,  and  the  solution  made  up 
to  250  c.c.  Fifty  cubic  centimetres  of  this  solution  are  then  with- 
drawn, and  gently  warmed  to  a  temperature  about  60°,  and  then 
titrated  with  deci-normal  permanganate.! 

*  A  method  for  estimating  copper  is  based  upon  a  similar  reaction.  The 
copper  is  first  reduced  either  to  cuprous  oxide  by  means  of  grape  sugar,  or  to 
the  metallic  state  by  means  of  pure  metallic  zinc,  and  the  reduced  product  is 
then  dissolved  in  a  mixture  of  ferric  chloride  and  hydrochloric  acid — 

Cu20  +  Fe2Cl6  +  2HC1  =  2CuCl2  +  H2O  +  2FeCl2 
Cu  +  Fe2Cl6  =  CuCl2  +  2FeQ2 

t  Other  metals  (e.g.  lead)  which  admit  of  complete  precipitation  in  the 
form  of  oxalates  or  basic  oxalates  (bismuth)  are  capable  of  being  estimated  on 


Oxidation  and  Reduction  Methods.  341 

(b]  Instead  of  determining  the  oxalic  acid  which  is  actually  in 
combination  with  the  calcium,  the  calcium  may  be  precipitated  by 
means  of  an  excess  of  a  standard  oxalic  acid  solution,  and  after 
removing  the  precipitate  by  nitration,  the  excess  of  oxalic  acid  is 
found  by  means  of  permanganate.  By  deducting  this  excess 
from  the  total  amount  used,  the  quantity  of  oxalic  acid  which  was 
precipitated  by  the  calcium  is  found,  and  from  this  the  percentage 
of  calcium  is  calculated. 

About  o'5  gram  of  calc-spar  is  weighed  out,  and  dissolved 
in  dilute  hydrochloric  acid  in  a  flask.  Twenty  cubic  centimetres 
of  normal  oxalic  acid  (63  grams  per  litre)  are  delivered  into  the 
flask,  and  a  slight  excess  of  ammonia  added.  The  mixture  is 
boiled  for  a  few  minutes,  and  then  allowed  to  cool,  after  which  it 
is  diluted  with  water  up  to  250  c.c.  The  liquid  is  then  filtered 
through  an  univetted  filter,  and  when  sufficient  has  passed  through, 
50  c.c.  of  the  perfectly  clear  filtrate  *  are  withdrawn  with  a  pipette 
and  transferred  to  a  small  flask,  in  which  the  solution  is  acidified 
with  sulphuric  acid,  heated  to  about  60°,  and  titrated  with  deci- 
normal  permanganate. 

Since  one- fifth  of  the  total  solution  was  used  for  the  titration, 
the  result  multiplied  by  5  will  give  the  excess  of  oxalic  acid  ;  and 
this  deducted  from  the  original  quantity  taken  will  give  the  amount 
which  was  precipitated  as  calcium  oxalate. 

Deci-normal  Potassium  Bichromate. 

(4-913  grams  per  litre.\] 

4-913  grams  of  pure  potassium  dichromate  (previously  dried  by 
being  gently  fused  in  a  porcelain  dish,  and  then  powdered  in  a  dry 
mortar)  are  exactly  weighed  out  and  dissolved  in  water  in  a  litre 
flask,  and  the  solution  made  up  to  1000  c.c.  with  water  at  15°. 

N 
One  litre  of  this  —  solution  contains  one-tenth  of  an  equivalent 

of  available  oxygen  in  grams,  i.e.  0*8  gram  ;  hence  I  c.c.  =  o'oooS 
gram  available  oxygen,  and  is  equivalent  to  0*0056  Fe,  or  to  the 
one-thousandth  of  the  equivalent  weight  of  any  substance  capable 
of  being  fully  oxidised  by  one  equivalent  of  oxygen. 

Unlike  potassium   permanganate,   the  dichromate  solution   is 

this  principle.  Thus  zinc  and  cadmium  oxalates  also  are  completely  preci- 
pitated by  oxalic  acid  on  the  addition  of  a  moderate  quantity  of  alcohol. 
Alkali  salts  must  be  absent,  as  these  produce  soluble  double  oxalates. 

*  Any  calcium  oxalate  passing  through  the  filter  will  be  subsequently 
decomposed  by  the  acid,  and  will  be  therefore  returned  as  excess  oxalic  acid, 
and  so  vitiate  the  analysis. 

f  See  p.  313. 


342  Volumetric  Analysis. 

stable,  and  as  it  has  no  action  upon  rubber,  it  may  be  used  in  a 
burette  which  is  closed  by  a  rubber  tube  and  pinchcock. 

Titration  of  Decinormal  Potassium  Bichromate. 

(a)  By  means  of  Ferrous  Sulphate — 

K2Cr207  +  6FeS04  +  7H2SO4  =  K2SO,  +  Cr0(SO4)3  +  3Fe.(SO4)3 

+  ?H20 

or,  since  the  dichromate  contains  3  atoms  of  available  oxygen,  the 
reaction  may  be  expressed — 

3O  +  6FeO  =  3Fe2O3 

For  the  purposes  of  the  titration,  the  ferrous  sulphate  is  prepared 
by  dissolving  pure  iron  in  dilute  sulphuric  acid,  with  exclusion  of 
air,  precisely  as  described  for  permanganate,  p.  333.  An  aliquot 
part  of  the  solution — say  50  c.c. — is  withdrawn  by  means  of  a 
pipette,  and  transferred  to  a  small  flask,  and  the  dichromate  solution 
gradually  added  from  a  burette. 

In  this  process,  the  end  of  the  reaction  is  ascertained  by  means 
of  a  freshly  made  and  dilute  solution  of  potassium  ferricyanide, 
used  as  an  outside  indicator.  A  number  of  drops  of  the  ferricyanide 
are  placed  about  upon  a  white  plate  or  tile,  and  from  time  to  time, 
during  the  addition  of  the  dichromate,  a  drop  of  the  mixture  is 
withdrawn  upon  a  glass  rod  and  brought  into  contact  with  one  of 
the  drops  of  the  indicator.  At  first  a  strong  blue  coloration  is 
produced,  but  as  the  amount  of  ferrous  salt  is  gradually  diminished 
by  the  addition  of  the  dichromate,  the  blue  becomes  less  and  less 
intense,  until  at  last  a  drop  of  the  liquid  so  tested  fails  to  give  any 
coloration.  At  this  point  the  whole  of  the  ferrous  salt  has  been 
oxidised,  and  the  reaction  is  therefore  complete.* 

(b)  By  means  of  Ferrous  Ammonium  Sulphate. — The  process 
is  carried  out  as  in  the  case  of  potassium  permanganate,  the  end 
reaction  being  determined  by  means  of  ferricyanide,  as  described 
above. 

Analyses  by  means  of  Potassium  Dichromate. 

In  a  number  of  instances,  potassium  dichromate  may  be  sub- 
stituted for  permanganate  in  volumetric  analysis.  This  is  the  case, 
for  example,  with  all  estimations  that  are  based  upon  the  oxidation 
of  ferrous  to  ferric  salts  ;  but  not  with  those  in  which  oxidation  of 
oxalic  acid  is  the  foundation  of  the  process. 

*  It  will  be  evident  that  it  is  absolutely  essential  to  the  success  of  this 
operation  that  the  ferricyanide  should  be  perfectly  free  from  ferrocyanide, 
otherwise  the  oxidised  iron  will  itself  give  rise  to  a  blue  coloration. 


Analyses  by  means  of  Potassium  Bichromate.     343 

The  precautions  as  to  the  presence  of  hydrochloric  acid,  which 
have  to  be  taken  when  using  potassium  permanganate,  do  not  apply 
in  the  case  of  the  dichromate ;  and  the  solution  to  be  titrated  may 
be  less  dilute  when  the  latter  reagent  is  employed. 

The  reducing  agent  to  be  used,  in  order  to  convert  the  iron  into 
the  ferrous  state  previous  to  titration  with  potassium  dichromate,  is 
either  an  alkaline  sulphite,  or  stannous  chloride.  Reduction  by 
means  of  metallic  zinc  (which  is  the  method  most  suitable  when 
permanganate  is  to  be  used,  see  p.  336)  is  not  so  well  adapted  for 
solutions  which  are  to  be  titrated  with  the  dichromate  solution,  as 
the  presence  of  the  dissolved  zinc  salt  interferes  with  the  reaction 
with  the  ferricyanide  indicator. 

(i)  Estimation  of  Iron  in  Iron  Ores.*— About  i  gram 
of  finely  powdered  and  dry  red  haematite  is  weighed  out  into  a 
flask  and  boiled  with  a  small  quantity  of  strong  hydrochloric  acid, 
diluted  with  about  half  its  own  volume  of  water,  until  the  whole  of 
the  iron  has  been  extracted,!  and  the  residue  is  free  from  dark- 
coloured  particles. 

Reduction  by  means  of  Stannous  Chloride. — The  solution  is  cooled 
and  diluted  up  to  100  c.c.  with  cold  water.  Twenty-five  cubic 
centimetres  of  this  solution  are  then  transferred  to  a  small  flask  by 
means  of  a  pipette,  and  heated  to  the  boiling-point  over  a  small 
flame.  A  clear  freshly  prepared  solution  of  stannous  chloride  is 
cautiously  added,  drop  by  drop,  until  the  yellow  colour  of  the  ferric 
chloride  is  just  discharged,  carefully  avoiding  unnecessary  excess — 

Fe2Cl6  +  SnCl2  =  SnO4  +  2FeCl2 

The  slight  excess  of  stannous  chloride  is  removed  from  the 
solution  (as  stannous  chloride  is  oxidised  by  potassium  dichromate, 
and  its  presence,  therefore,  would  vitiate  the  subsequent  titration), 
by  adding  a  few  drops  of  a  solution  of  mercuric  chloride,  whereby 

*  In  the  example  here  given,  the  total  iron  is  alone  estimated.  When  the 
amount  of  ferrous  and  ferric  oxide  present  is  to  be  separately  determined,  the 
procedure  is  the  same  as  that  described  on  p.  335  ;  the  reduction  of  the  iron 
in  the  second  portion  (in  which  the  total  iron  is  estimated)  is  accomplished 
either  by  stannous  chloride  or  alkaline  sulphite,  and  the  solutions  titrated  with 
dichromate. 

f  In  cases  where  the  iron  is  difficult  to  extract  with  hydrochloric  acid  alone, 
a  few  small  crystals  of  potassium  chlorate  may  be  added,  after  which  the 
solution  must  be  evaporated  to  dryness,  and  the  residue  dissolved  in  hydro- 
chloric acid.  Iron  ores  (such  as  the  titaniferous  ores)  from  which  the  iron 
cannot  be  wholly  extracted  by  acids,  are  mixed  in  a  state  of  fine  powder  with 
from  6  to  10  times  their  weight  of  hydrogen  potassium  sulphate  (bisulphate  of 
potash],  and  gently  fused  in  a  platinum  dish  for  from  20  to  30  minutes.  The 
"  melt"  is  extracted  with  dilute  hydrochloric  acid. 


344  Volumetric  Analysis. 

mercur0#.y  chloride  is  precipitated  (which  is  without  action  upon 
the  dichromate),  and  stannic  chloride  produced  * — 

SnCl2  +  2HgCl2  =  Hg2Cl2  +  SnCl4 

The  solution  is  now  titrated  with  the  deci-normal  dichromate 
without  unnecessary  exposure  to  the  air.  The  volume  of  the  di- 
chromate used,  multiplied  by  4  (as  25  c.c.  out  of  the  100  c.c.  were 
used  for  the  titration),  gives  the  volume  required  to  oxidise  the  iron 
in  the  original  weight  of  ore  taken. 

Reduction  by  means  of  Alkaline  Sulphite. — The  solution  of  the 
ore  obtained  by  boiling  with  hydrochloric  acid  is  slightly  diluted, 
and,  if  necessary,  filtered  through  a  small  filter  into  a  flask,  the 
filter  and  residue  being  thoroughly  washed.  Ammonia  is  added  to 
the  liquid  until  slightly  alkaline,  as  shown  by  the  formation  of  a 
faint  permanent  precipitate  of  ferric  hydroxide.  One  or  two  drops 
of  dilute  hydrochloric  acid  are  then  added,  in  order  to  just  re-dissolve 
this  precipitate. 

Ten  cubic  centimetres  of  a  strong  solution  of  hydrogen  ammonium 
sulphite  f  (ammonium  bisulphite),  H(NH4)SO3,  are  then  added, 
and  thoroughly  mixed  by  shaking,  and  the  solution  gradually  heated 
to  boiling.  Ten  cubic  centimetres  of  strong  sulphuric  acid  are 
mixed  with  an  equal  volume  of  water,  and  this  mixture  is  added  to 
the  heated  liquid,  and  the  boiling  continued  for  15  or  20  minutes 
to  expel  the  whole  of  the  sulphur  dioxide.  The  solution  is  then 
cooled  quickly  and  diluted  to  100  c.c.  with  cold  water.  Twenty-five 
cubic  centimetres  of  this  are  then  titrated  with  the  deci-normal 
dichromate. 

(2)  Estimation  of  Chromium  in  Chrome  Iron  Ore. 

By  the  exact  converse  of  the  foregoing  process,  chromium  (in 
the  condition  of  chromate)  is  most  readily  estimated.  In  the  case 
of  chrome  iron  ores  (as  also  with  ferro-chromium),  the  chromium 
is  converted  into  sodium  chromate  by  fusion  with  sodium  peroxide. 
The  chromate  is  then  reduced  by  means  of  a  measured  or  weighed 
excess  of  a  ferrous  salt,  and  the  excess  of  the  latter  then  estimated 
by  titration  with  deci-normal  dichromate  solution. 

*  If  no  precipitate  is  produced  on  the  addition  of  the  mercuric  chloride,  it 
shows  that  an  insufficient  quantity  of  stannous  chloride  was  added.  Too 
copious  a  precipitate,  due  to  too  great  an  excess  of  the  tin  solution,  will 
interfere  with  the  titration. 

f  That  is  to  say,  about  i  c.c.  for  each  o'i  gram  of  ore  which  was  dissolved. 
The  bisulphite  may  be  prepared  by  passing  a  gentle  stream  of  sulphur  dioxide 
through  a  small  quantity  of  strong  ammonia  solution,  diluted  with  half  its 
volume  of  water.  The  ammonia  may  be  contained  in  a  boiling-tube,  which  is 
kept  cool  by  being  immersed  in  a  beaker  of  cold  water.  The  gas  is  bubbled 
through  it  until  the  liquid  smells  strongly  of  sulphur  dioxide,  and  until  the 
crystals  of  the  normal  sulphite  which  first  form  are  wholly  re-dissolved. 


Analyses  by  means  of  Potassium  Bichromate.     345 

About  0*5  gram  of  \htfinely  powdered  ore  *  is  weighed  out  into 
a  nickle  crucible  f  and  intimately  mixed  with  3  to  4  grams  of  sodium 
peroxide.  The  crucible  is  then  heated  by  means  of  a  Bunsen  flame, 
and  the  mixture  kept  in  a  state  of  fusion  for  5  to  10  minutes  It  is 
then  allowed  to  partially  cool,  until  a  crust  is  formed  upon  the 
surface  of  the  fused  mass,  when  I  gram  of  sodium  peroxide  is  added, 
and  the  materials  again  melted  and  kept  in  a  state  of  fusion  for 
5  minutes  longer. 

After  cooling,  the  crucible  is  transferred  to  a  porcelain  beaker 
(or,  in  the  absence  of  this,  a  porcelain  dish),  and  the  melt  exhausted 
with  hot  water,  the  beaker  or  dish  being  covered  with  a  clock-glass. 
The  solution  is  then  boiled  for  10  minutes,  in  order  to  decompose 
completely  the  whole  of  the  remaining  sodium  peroxide.^  [Should 
the  solution  first  obtained  by  dissolving  the  melt  be  purple  in  colour, 
due  to  sodium  ferrate,  a  little  more  sodium  peroxide  should  be  added 
before  boiling  the  liquid.] 

The  iron  present  in  the  ore  (the  manganese,  in  the  case  of  ferro- 
chromium),  and  the  nickel  derived  from  the  crucible,  are  left  in 
the  form  of  oxides,  which  are  separated  by  filtering  and  carefully 
washing  the  residue  upon  the  filter.  The  yellow  solution  containing 
the  chromium,  as  sodium  chromate,  is  acidulated  with  sulphuric 
acid,  and  an  excess  of  ferrous  ammonium  sulphate  added.  This  is 
most  conveniently  done  by  adding  a  weighed  quantity  of  the  solid 
salt.  A  weighing-bottle,  filled  with  the  salt,  is  first  weighed,  and 
then  small  quantities  are  carefully  taken  out  upon  a  small  clean 
spatula  (holding  the  bottle  over  the  beaker  containing  the  chromate 
solution,  lest  any  particles  should  fall),  and  added  to  the  liquid 
until  the  chromium  is  entirely  reduced  (the  solution  turning  green) 

*  Too  much  stress  cannot  be  laid  upon  the  importance  of  reducing  the  ore  to 
the  finest  possible  powder,  as  the  success  of  the  "  fusion  "  hangs  entirely  upon 
this  point.  The  ore  should  be  crushed  in  a  steel  mortar  until  the  powder  is 
fine  enough  to  pass  through  a  linen  bag  ;  it  is  then  ground  in  an  agate  mortar 
with  the  addition  of  a  little  water,  until  it  is  reduced  loan  impalpable  powder. — 
Rideal  and  Rosenblum  (the  authors  of  the  method),  Journal  of  the  Society 
of  C hem.  Industry,  December  31,  1895. 

f  Platinum  and  silver  vessels  are  attacked  by  the  fused  peroxide.  Nickel  is 
also  attacked  to  some  extent,  but  the  dissolved  nickel  is  subsequently  removed. 

£  It  is  essential  to  the  success  of  the  process  that  the  whole  of  the  sodium 
peroxide  be  decomposed  at  this  stage,  otherwise,  when  the  liquid  is  subsequently 
acidified,  hydrogen  peroxide  will  be  generated,  and  this,  reacting  upon  the 
chromate  in  solution,  causes  the  reduction  of  a  portion  of  the  chromium.  The 
reducing  action  of  hydrogen  peroxide  upon  chromic  acid  is  similar  in  its 
character  to  that  which  it  exerts  upon  other  highly  oxidised  compounds  (e.g. 
peroxides,  permanganates,  etc.),  whereby  oxygen  is  withdrawn  from  both  com- 
pounds, and  is  liberated  as  free  gas.  In  the  case  of  chromium,  however,  an 
intermediate  stage  occurs,  in  which  the  blue  compound,  supposed  to  be  per- 
chromic  acid,  is  first  formed. 


346  Volumetric  Analysis. 

and  the  ferrous  salt  is  present  in  slight  excess.  That  excess  of  iron 
has  been  introduced  is  ascertained  by  bringing  a  drop  of  the  liquid 
(by  means  of  a  glass  rod)  into  contact  with  a  drop  of  a  fresh  and 
dilute  solution  of  potassium  ferricyanide  upon  a  white  plate,  when 
a  blue  colour  is  obtained.  The  stopper  is  replaced  in  the  weighing- 
bottle,  and  the  latter  set  aside  to  be  re-weighed  after  the  titration 
has  been  completed. 

The  excess  of  ferrous  salt  present  is  now  determined  by  means 
of  deci-normal  potassium  dichromate,  which  is  added  until  a  drop 
of  the  liquid  no  longer  gives  a  blue  colour  with  potassium 
ferricyanide. 

The  weighing-bottle  is  now  re-weighed,  the  difference  being 
the  total  weight  of  the  ferrous  salt  used.  From  this  is  deducted 
the  weight  of  ferrous  salt  equivalent  to  the  volume  of  deci-normal 
dichromate  employed,  i.e.  the  excess  of  ferrous  salt,  which  gives 
the  weight  of  ferrous  salt  oxidised  by  the  chromate  derived  from 
the  ore.  From  the  equation  2CrO3  +  6FeO  =  Cr2O3  +  3Fe2O3, 
it  will  be  seen  that  I  atom  of  Cr  (52*4)  is  capable  of  oxidising  3 
atoms  of  Fe  (56  x  3  =  168). 

The  formula-weight  of  ferrous  ammonium  sulphate  being  exactly 
seven  times  the  weight  of  the  iron  it  contains  (see  p.  334),  the 
weight  of  this  salt  which  is  equivalent  to  52-4  parts  of  chromium  is 
1176  parts.  From  this  proportion  the  percentage  of  chromium  in 
the  ore  is  readily  calculated. 

Iodine. 

Free  iodine,  in  the  presence  of  water  and  certain  oxidisible 
substances,  is  able  to  take  the  hydrogen  from  water  and  liberate 
the  oxygen.  The  oxygen  thus  eliminated  oxidises  the  oxidisible 
compound  present,  and  therefore  the  iodine  acts  as  an  indirect 
oxidising  agent,  127  parts  of  iodine  being  equivalent  to  8  parts  of 
oxygen. 

Thus,  when  a  solution  of  iodine  is  added  to  dilute  sulphurous 
acid,  the  latter  compound  is  oxidised  to  sulphuric  acid — 
H2S03  +  I2  +  H20  -  H2S04  +  2HI 

Similarly,  the  lower  oxides  of  arsenic  and  of  antimony  (or  the 
salts  of  these  oxides)  are  oxidised  by  iodine  in  the  presence  of 
water  ;  thus  with  sodium  arsenite— 

Na3AsO3  +  I2  +  H2O  =  Na3AsO4  +  2HI 

Therefore  a  standard  solution  of  iodine  may  be  used  to  estimate 


Oxidation  Reactions  by  means  of  Iodine,        347 

sulphur  dioxide  in   solution,   and   also   for  the   determination  of 
arsenic  and  antimony. 

When  iodine  is  added  to  a  dilute  solution  of  sulphuretted 
hydrogen,  a  definite  reaction  takes  place,  which  is  usually  expressed 
by  the  equation  * — 


Hence  a  solution  of  iodine  may  be  employed  for  the  direct 
determination  of  sulphuretted  hydrogen  in  solution.  Of  all  the 
reactions  for  which  a  solution  of  iodine  is  employed,  the  most 
useful  and  important  in  volumetric  analysis  is  that  with  thio- 
sulphuric  acid,  or  sodium  thiosulphate,  which  becomes  converted 
into  the  sodium  salt  of  tetrathionic  acid,  while  hydriodic  acid  (or 
the  sodium  salt)  is  produced  ;  thus — 

2Na2S203  +  I2  =  2NaI  +  Na2S4OGt 

In  volumetric  processes  in  which  iodine  is  employed,  the  end 
reaction  is  indicated  by  means  of  starch.  When  brought  in  contact 
with  free  iodine,  starch  produces  a  deep  blue  coloration,  and  the 
formation  of  this  blue  colour  constitutes  a  delicate  indication  of 
the  smallest  excess  of  iodine  over  and  above  that  which  is  required 
for  the  oxidising  reaction  in  hand. 

The  blue-coloured  substance  is  usually  regarded  as  a  compound 
of  starch  and  iodine,  although  unstable  in  character  and  of  un- 
known composition ;  the  idea  of  combination  between  the  substances 
is  also  conveyed  by  the  name  iodide  of  starch,  by  which  it  is  called. 
Towards  sodium  thiosulphate,  however,  in  common  with  many 
other  reagents,  it  behaves  exactly  as  free  iodine  :  if  a  solution  of 

*  In  this  equation  the  oxidising  character  of  the  process  is  obscured  by 
the  omission  from  it  of  the  water  which  is  necessary  for  its  progress.  At  the 
ordinary  temperature,  hydrogen  sulphide  and  iodine  do  not  interact  in  the 
absence  of  water.  One  function  of  the  water  is,  by  dissolving  the  hydriodic 
acid,  to  furnish  the  necessary  heat  required  for  the  combination  of  the  hydrogen 
and  iodine ;  while  another,  no  doubt,  is  to  supply  the  intermediate  step  in  the 
process  whereby  oxygen  is  available  for  the  oxidation  of  the  sulphuretted 
hydrogen,  thus — 

|I2  +  H2:;0  +  HJS  =  2HI  +  H20  +  S 


f  This  reaction,  just  as  in  the  case  of  sulphuretted  hydrogen,  is  in  reality 
one  of  indirect  oxidation  through  the  intervention  of  water.  The  mechanism 
of  the  reaction  will  be  understood  by  the  following  dissected  equations — 


;!2  +  H2:jO  +  NaJNaS2O3  =  *Hl  +  ONaz  +  Na2S<°« 
2HI  -f-  ONa2  =  2NaI  +  H2O 


348  Volumetric  Analysis. 

the  thiosulphate  be  added  to  it,  the  iodine  at  once  attacks  the 
thiosulphate  according  to  the  above  reaction,  just  as  though  there 
were  no  starch  present ;  and  the  instant  the  whole  of  the  iodine 
is  thus  used  up,  the  blue  colour  of  the  liquid  disappears. 

It  is  on  this  account  that  a  solution  of  sodium  thiosulphate  used 
in  conjunction  with  the  iodine  solution  constitutes  so  important 
a  volumetric  reagent.  Thus,  all  substances  which  can  be  made 
to  yield  chlorine  either  directly  (as  from  hypochlorites  or  chlorates) 
or  indirectly  (as  by  the  action  of  hydrochloric  acid  upon  peroxides, 
chromates,  etc.)  are  capable  of  being  estimated  by  these  reagents. 
The  liberated  chlorine  is  passed  into  potassium  iodide,  and  there 
displaces  its  equivalent  of  iodine,  and  the  exact  amount  of  iodine 
can  then  be  determined  by  means  of  sodium  thiosulphate. 

Deci-normal  Iodine  Solution. 

(127  grams  of  iodine  per  litre?) 

127  grams  of  the  purest  re-sublimed  iodine,  in  a  state  of 
powder,  are  exactly  weighed  out  into  a  litre  flask.  About  20  grams 
of  potassium  iodide  (free  from  iodate)  dissolved  in  about  200  c.c. 
of  water  are  added,  and  the  mixture  gently  shaken  until  the  iodine 
is  wholly  dissolved.  The  solution  is  then  diluted  up  to  the  litre 
with  water  at  15°,  and  preserved  in  a  well- fitting  stoppered  bottle 
in  the  dark. 

If  strictly  deci-normal,  i  c.c.  will  contain  0-0127  grams  of  iodine, 
equivalent  to  0*0008  gram  oxygen  or  to  0*00355  gram  of  chlorine. 

Deci-normal  Sodium  Thiosulphate. 

(24*8  grams  ofNa.2S2O3,$H2O  per  litre.} 

24*8  grams  of  the  purified  salt,  which  has  been  carefully 
powdered,*  are  weighed  out  into  a  litre  flask,  and  dissolved  in 
water.  The  solution  is  then  diluted  up  to  icoo  c.c.  with  water  at 
15°.  If  strictly  deci-normal,  i  c.c.  of  the  solution  will  be  exactly 
oxidised  by  i  c.c.  of  the  iodine  solution. 

Titration  of  Deci-normal  Iodine,  by  means  of  sodium 
thiosulphate. 

*  When  powdering  a  salt  which  contains  a  large  number  of  molecules  of 
water  of  crystallisation,  such  as  sodium  thiosulphate,  there  is  always  a  risk  of 
a  slight  loss  of  water.  This  is  due  to  the  fact  that  the  pressure  produced  by 
the  blow  of  the  pestle  causes  the  momentary  partial  liquefaction  of  portions  of 
the  salt.  This  explains  the  fact  that,  although  perfectly  free  from  extraneous 
moisture,  such  a  salt  appears  to  become  damp,  and  ' '  clogs  "  upon  the  pestle 
during  the  process  of  powdering.  While  in  this  condition  the  salt  may  lose 
traces  of  its  water  of  crystallisation  by  evaporation. 


Iodine  and  Sodium   ThiosulpJiate.  349 

Twenty- five  cubic  centimetres  of  the  deci-normal  thiosulphate 
solution  are  transferred  by  means  of  a  pipette  to  a  beaker,  diluted 
with  a  little  water,  and  a  few  drops  of  a  clear  solution  of  starch 
added.*  The  iodine  solution  is  then  run  in  from  a  stoppered 
burette,  until  a  single  drop  gives  a  permanent  blue  coloration. 

The  titration  may  equally  as  well  be  made  in  the  opposite 
order.  Twenty-five  cubic  centimetres  of  the  iodine  solution  are 
transferred  to  a  beaker,  diluted  with  about  an  equal  volume  of 
water,  and  the  starch  solution  added.  The  thiosulphate  is  then 
delivered  from  a  burette  until  the  blue  colour  is  just  discharged. 
As  a  rule,  when  the  titration  is  made  in  this  way,  it  is  an  advantage 
not  to  introduce  the  starch  until  towards  the  end  of  the  reaction, 
especially  if  the  amount  of  iodine  is  considerable.  The  thio- 
sulphate is  added  until  the  brown  coloiyr  of  the  iodine  solution 
begins  to  perceptibly  pale,  and  the  liquid  assumes  a  straw  colour. 
One  or  two  drops  of  starch  are  then  introduced,  causing  the  blue 
colour,  and  the  thiosulphate  delivered  drop  by  drop  until  the 
colour  is  discharged. 

The  result  of  this  titration  will  show  the  exact  relation  in  which 
these  two  solutions  stand  to  each  other.  If  they  bear  the  correct 
relative  values,  then  25  c.c.  of  the  one  will  require  25  c.c.  of  the 
other  in  order  to  exactly  complete  the  reaction  between  them.  If, 
on  the  other  hand,  one  is  found  to  be  proportionately  a  little  stronger 
than  the  other,  the  value  of  the  one  in  terms  of  the  other  must  be 
notified,  preferably  on  the  labels  of  the  bottles.  Thus,  suppose  the 
25  c.c.  of  iodine  solution  required  24-9  c.c.  of  the  thiosulphate  solu- 
tion, then  i  c.c.  of  the  latter  equals  —  5  or  1-004  c.c.  of  the  iodine  ; 

24'9 
and  I  c.c.  of  the  iodine  is  equivalent  to  0-996  c.c.  of  thiosulphate. 

Although  these  solutions  might  be  found  to  bear  the  exact 
relative  value  towards  each  other,  so  that  equal  volumes  are 
chemically  equivalent,  it  is  obviously  possible  that  they  might  still 
not  be  strictly  deci-normal  solutions.  It  is  possible,  although  per- 
haps not  highly  probable,  that  an  error  to  exactly  the  same  extent 
might  have  arisen  in  the  preparation  of  each  ;  so  that  each  of  them 

*  The  starch  solution  is  prepared  by  mixing  about  i  or  2  grams  of  white 
potato-starch  in  powder  with  5  or  6  c.c.  of  cold  water  in  the  bottom  of  a 
moderately  large  beaker.  A  considerable  quantity  of  boiling  water  is  then 
poured  upon  it.  As  the  hot  water  is  being  added,  the  opaque  white  appearance 
which  the  mixture  presents  at  first,  changes  almost  suddenly  to  that  of  a  semi- 
translucent  gelatinous  substance.  At  this  point  the  addition  of  the  boiling 
water  is  stopped,  and  the  beaker  nearly  filled  up  with  cold  water.  It  is 
allowed  to  settle,  and  the  clean  liquid  poured  off  for  use.  This  starch  solution 
should  nof  be  kept  for  more  than  one  day,  or  at  the  outside  two  days. 

2    A 


3$o  Volumetric  Analysis. 

should  be  either  too  strong  or  too  dilute  to  the  s.ime  degree. 
It  is  necessary,  therefore,  to  make  an  independent  titration  of  one 
of  them  to  serve.  as  a  check.  Either  of  the  following  methods  can 
be  used. 

(a)  Titration  of  Deri  -normal  Iodine  by  means  of  Arsenious 
Oxide  — 

As406  4-  4l2  +  4H2O  =  SHI  +  2As2O5 

That  is  to  say,  127  parts  of  iodine  will  oxidise  49*5  parts  of 
arsenious  oxide. 

4'95  grams  of  finely  powdered  arsenious  oxide,  which  has  been 
purified  by  resublimation,  are  exactly  weighed  out  into  a  litre  flask, 
and  about  500  c.c.  of  water  added.  Thirty  grams  of  pure  sodium 
bicarbonate  are  then  introduced,  and  the  mixture  gently  warmed 
upon  a  steam-bath  to  a  temperature  about  60°,  and  continually 
shaken  until  the  arsenious  oxide  is  wholly  dissolved.*  The  solu- 
tion is  then  cooled  and  made  up  to  a  litre  with  water  at  1  5°.  This 
solution,  containing  one-tenth  of  the  equivalent  of  arsenious  oxide, 
is  of  deci-normal  strength,  i  c.c.  being  equivalent  to  0*0127  gram 
of  iodine,  and  therefore  equal  to  i  c.c.  of  deci-normal  iodine  solu- 
tion. 

Twenty-five  cubic  centimetres  of  this  solution  are  transferred  to 
a  small  beaker,  a  few  drops  of  clear  starch  solution  added,  and  the 
deci-normal  iodine  solution  delivered  into  it  from  a  burette  until 
the  blue  coloration  is  permanent.  From  the  volume  of  the  iodine 
solution  required  to  completely  oxidise  the  arsenious  oxide  in  25  c.c. 
of  the  solution,  the  exact  strength  of  the  iodine  solution  is  obtained. 
If  it  is  found  to  be  strictly  deci-normal,  then  obviously  the  thio- 
sulphate  solution  which  has  the  same  relative  strength  is  also 
exactly  deci-normal. 

(b~)  Titration  of  Deci-normal  Thiosulphate  Solution  by  means  of 
Potassium  Dichromate.  —  When  a  solution  of  potassium  dichromate, 

*  An  excess  of  the  alkali  is  necessary  in  order  to  neutralise  the  hydriodic 
acid  which  is  afterwards  produced  during  the  titration,  the  necessary  alkali 
for  this  purpose  being  supplied  by  sodium  bicarbonate,  instead  of  either  the 
normal  carbonate  or  caustic  alkali.  As  already  stated,  the  iodine  in  this  blue 
compound  behaves  towards  many  reagents  much  a.?,  free  iodine  would,  and  free 
iodine  is  dissolved  by  caustic  alkali,  and  also  (but  less  easily)  by  the  normal 
carbonates  of  the  alkalies  ;  thus  — 


6NaHO  +  3lz  =  5NaI  +  NaIO3  +  3H2O 
6Na2CO3  +  sI2  +  6H20  =  6NaHCO3  +  sNal  +  NaIO3 


Hence  if  either  of  these  alkaline  reagents  be  added  to  the  blue  starch  iodide, 
the  blue  colour  will  be  destroyed  (see  note  on  p.  355). 


Iodine  and  Sodium  Thiosulphate.  351 

acidified  with  hydrochloric  acid,  is  mixed  with  potassium  iodide, 
and  the  mixture  gently  warmed,  iodine  is  liberated  ;  thus  — 

K2Cr207  +  I4HC1  +  6KI  =  3!,  +  Cr2Cl6  +  8KC1  +  ;H2O 

The  amount  of  iodine  so  liberated  is  the  equivalent  of  the  available 
oxygen  in  the  dichromate,  three  atoms  of  available  oxygen  being 
equivalent  to  six  atoms  of  iodine,  i.e.  8  parts  of  oxygen  liberate  1  27 
parts  of  iodine. 

If  a  standard  solution  of  the  dichromate  be  used  in  the  reaction, 
the  available  oxygen  is  a  known  quantity,  and  therefore  the  amount 
of  iodine  it  can  liberate  is  also  a  known  quantity  ;  so  that  if  this 
iodine  be  titrated  with  sodium  thiosulphate,  the  strength  of  the 
latter  can  be  determined  — 

Thus  i  c.c.  deci-normal  dichromate  solution  contains  0*0008  gram 

available  oxygen  ; 
therefore  I  c.c.  deci-normal  dichromate  solution  liberates  0*0127  gram 

iodine. 

i  c.c.  deci-normal  thiosulphate  oxidises  0*0127  gram  iodine  ; 
therefore  i  c.c.  deci-normal  dichromate  is  equivalent  to  i  c.c.  deci- 

normal  thiosulphate. 

Twenty-five  cubic  centimetres  of  deci-normal  dichromate  solution 
are  placed  in  a  stoppered  bottle,  and  about  2  grams  of  potassium 
iodide  together  with  5  or  6  c.c.  of  strong  hydrochloric  acid  are  added. 
The  mixture  is  thoroughly  shaken,  and  gently  heated  by  standing 
the  closed  bottle  in  warm  water  for  some  time.  The  bottle  with 
its  contents  is  then  cooled,  the  stopper  withdrawn  and  rinsed  into 
the  bottle  with  water,  and  the  liberated  iodine  titrated  with  the 
thiosulphate  solution  ;  the  starch  solution  being  added  when  the 
reddish  colour  of  the  iodine  solution  begins  to  fade.  The  dis- 
appearance of  the  blue  colour  of  the  iodide  of  starch  leaves  the 
solution  not  colourless,  but  pale  green  (due  to  the  chromic  chloride), 
hence  a  little  care  is  necessary  in  judging  the  end  reaction.*  If 
the  deci-normal  thiosulphate  is  of  exact  strength,  25  c.c.  will  be 
required  for  the  titration. 

*  Deci-normal  permanganate,  acidified  with  dilute  sulphuric  acid,  may  be 
substituted  for  the  dichromate  and  hydrochloric  acid,  the  operation  being  con- 
ducted in  the  same  manner.  In  this  case,  when  the  blue  colour  of  the  starch 
iodide  is  discharged  by  the  thiosulphate,  the  liquid  is  left  colourless.  The  5 
atoms  of  available  oxygen  in  permanganate  liberate  10  atoms  of  iodine  ;  thus— 


2KMnO4  +  8H2S04  +  loKI  =  5I2  +  6K2SO4  +  2MnSO4  +  8H2O 


352  Volumetric  Analysis. 

Estimations  by  means  of  Deci-normal  Iodine. 

(1)  Antimony  in  Antimonious  Oxide. 

Sb2O3  +  2H2O  +  2l2  =  2HI  +  Sb2O5 

About  I  gram  of  antimonious  oxide  is  weighed  out  into  a  ^-litre 
flask.  Two  grams  of  tartaric  acid  dissolved  in  about  25  c.c.  of 
water  are  then  added,  and  the  mixture  gently  shaken  until  the 
antimonious  oxide  is  dissolved.  Sodium  carbonate  (the  normal 
salt)  is  added  until  the  solution  is  just  neutral,  and  the  liquid  is 
diluted  with  water  to  250  c.c.  Fifty  cubic  centimetres  of  this  solu- 
tion are  withdrawn  with  a  pipette  and  transferred  to  a  beaker,  and 
from  20  to  30  c.c.  of  a  cold  saturated  solution  of  pure  sodium 
bicarbonate*  added,  together  with  a  few  drops  of  starch.  The 
deci-normal  iodine  solution  is  then  run  in  from  a  burette  until  a 
permanent  blue  colour  is  produced.  A  duplicate  titration  should 
be  made  in  a  second  portion  of  the  solution. 

EXAMPLE. — Fifty  cubic  centimetres  of  the  solution  required 
27*5  c.c.  of  the  iodine  solution  ; 

therefore  250  c.c.  (containing  i  gram  of  the  original  oxide)  would 

require  27-5  X  5  =  137-5 
i  c.c.  deci-normal  iodine  =  0-0127  gram  I  =  o'oo6o  gram  Sb 

=  0-0072  gram  Sb2O3 

therefore   I37'5  x  0*0060  =  0-8250  gram  Sb  in  i  gram  of  the  oxide 
0-8250  x  loo  =  82.5O   percentage   Sb,  or   99  0  per- 
centage of  Sb2O3 

(2)  Arsenic  in  Arsenious  Compounds.— The  process  is 
carried  out  as  already  described  on  p.  350. 

(3)  Tin,  in  the  Condition  of  Stannous  Chloride,  may 

also  be  estimated  by  means  of  iodine,  but  in  this  case  the  principle 
involved  is  different  to  that  which  obtains  with  either  arsenic  or 
antimony.  Stannous  chloride  unites  directly  with  iodine,  forming 
the  compound  SnCl2J2 — 

SnCl2  +  I2  =  SnCl2I2 

hence  two  atoms  of  iodine  become  the  measure  of  one  atom  of  tin, 
or  127  parts  of  iodine  are  equivalent  to  59  parts  of  tin.  Tin  cannot, 
however,  be  determined  by  this  reaction  in  the  presence  of  lead  ; 
because,  although  lead  chloride  alone  is  not  acted  upon  by  free 
iodine,  it  is  acted  upon  by  the  stannic  chlor-iodide  with  the  pre- 
cipitation of  yellow  lead  iodide,  which  completely  obscures  the  end 
reaction.  In  tin  alloys,  therefore,  which  contain  lead,  the  latter 
metal  must  be  first  removed. 

*  See  note  on  p.  350. 


Valuation  of  Bleaching  Powder.  353 

Estimations  by  means  of  Iodine  and  Thiosulphate. 

(1)  Sulphur  Dioxide  in  Sodium  Sulphite,  Na2SO3,7H2O — 

Na2S03  +  I2  +  H20  =  2HI  +  Na2SO4 

That  is  to  say,  1 26  parts  of  the  crystallised  salt  are  oxidised  by 

N 
127  parts  of  iodine;  therefore  i  c.c.  —  iodine  (containing  0*0127 

gram  I)  is  equivalent  to  0-0126  gram  Na2SO3,7H2O,  or  0-0032  gram 
S02. 

About  i  gram  of  the  powdered  salt  is  weighed  out  into  a  flask, 
and  loo  c.c.  of  the  iodine  solution  added.  This  volume  ensures  an 
excess  of  iodine,  so  long  as  the  weight  of  salt  taken  has  not  exceeded 
1-25  grams,*  and  the  excess  will  be  manifest  by  the  colour  of  the 
liquid. 

The  excess  of  iodine  is  now  titrated  by  means  of  the  deci-normal 
thiosulphate  solution,  added  from  a  burette.  As  soon  as  the  brown 
colour  of  the  iodine  begins  to  pale  to  a  yellow,  starch  is  added,  and 
then  the  addition  of  the  thiosulphate  is  continued  until  the  blue 
colour  is  discharged.  The  excess  of  iodine  thus  found,  deducted 
from  the  volume  originally  taken  (100  c.c.  in  this  case),  gives  the 
amount  of  iodine  which  has  been  used  in  oxidising  the  sodium  sul- 
phite in  accordance  with  the  above  reaction,  and  from  this  the 
percentage  of  sulphur  dioxide  can  be  calculated.! 

(2)  Available  Chlorine  in   Bleaching   Powder.  —  The 
active  ingredient  in  bleaching  powder  is  the  compound  expressed 
by  the  formula  Ca(OCl)Cl.     Bleaching  powder,  however,  besides 
containing   some   unacted-upon  calcium  hydroxide,  usually    con- 
tains  also   more   or  less   chlorinated    compounds   (such   as   cal- 
cium chlorate,  for  example),  which  are  of  no  use  for  the  purposes 
to  which  the  bleaching  powder  is  put.      The  chlorine  which  is 
present  as  calcium  chloro-hypochlorite,  Ca(OCl)Cl,  is  known  as  the* 
available  chlorine.     This  chlorine  is  evolved  when  the  bleaching 
powder  is  acted  upon  by  dilute  acids  ;  thus — 

Ca(OCl)Cl  +  H2SO4  =  H2O  4-  CaSO4  +  C12 
Ca(OCl)Cl  +  2HC1  =  H2O  +  CaCl2  +  C12 

Hence,  if  bleaching  powder  consisted  of  the  pure  compound, 

*  In  cases  where  the  composition  of  the  sulphite  is  unknown,  a  preliminary 
titration  must  be  made,  using  a  considerable  quantity  of  iodine,  so  as  to  gain 
an  approximate  idea  of  the  amount  required. 

f  This  method  may  also  be  applied  for  the  estimation  of  sulphuretted 
hydrogen  in  solution  (as  in  hepatic  waters),  and  also  of  aqueous  solutions  of 
chlorine  and  bromine. 


354  Volumetric  Analysis. 

having  the  formula  Ca(OCl)Cl,  it  would  contain  55  per  cent,  of 
available  chlorine  ;  but  in  reality  this  is  never  the  case>  and  the 
best  samples  of  commercial  bleaching  powder  seldom  exceed  about 
36  per  cent.,  while  in  others  it  falls  considerably  below  this. 

About  10  grams  of  bleaching  powder  are  weighed  out  into  a 
porcelain  mortar,  and  rubbed  down  with  a  small  quantity  of  water 
until  the  mixture  has  the  consistency  of  thin  cream.  The  heavier 
particles  are  allowed  a  moment  or  two  to  settle,  and  the 
milky  liquid  is  poured  off  into  a  litre  flask.  The  residue  is  then 
ground  down  with  a  little  more  water,  and  the  process  repeated 
until  the  last  traces  have  been  transferred  to  the  flask.  The  mix- 
ture is  then  diluted  with  water  up  to  1000  c.c.,  and  thoroughly 
shaken. 

Twenty-five  cubic  centimetres  of  the  milky  fluid  are  transferred 
by  means  of  a  pipette  to  a  small  beaker,  the  contents  of  the  flask 
being  briskly  shaken  up  immediately  before  the  withdrawal  of  the 
sample.  Excess  of  potassium  iodide  is  added  by  introducing  10 
c.c.  of  a  solution  containing  5  grams  KI  in  100  c.c.  water.*  The 
mixture  is  then  acidified  with  acetic  acid.f  This  liberates  the 
chlorine  from  the  bleaching  powder,  and  that,  in  the  presence  of 
the  excess  of  potassium  iodide,  sets  free  an  equivalent  quantity  of 
iodine — 

C12  +  2KI  =  2KC1  +  I2 

The  amount  of  the  liberated  iodine  is  then  estimated  by  the  addition 
of  deci-normal  sodium  thiosulphate,  after  the  addition  of  starch. 

N 
i  c.c.  —  thiosulphate  =  0*0127  gram  I  =  0-00355  gram  Cl 

Therefore  the  volume  of  thiosulphate  required  to  discharge  the 
blue  colour  of  the  starch  iodide,  is  a  measure  of  the  chlorine  which 
was  liberated  from  the  bleaching  powder  contained  in  25  c.c.  of  the 
'original  mixture.  From  this  the  percentage  of  available  chlorine  is 
calculated. 

Valuation  of  Bleaching  Powder.  A  It er native  Method  by  means 
of  Arsenions  Oxide.\ — Instead  of  estimating  the  available  chlorine 

*  It  is  convenient  to  remember  that  a  suitable  excess  of  potassium  iodide 
will  be  ensured  if  twice  the  weight  of  potassium  iodide  is  added  as  there  is  of 
bleaching  powder  present.  In  the  above  example,  25  c.c.  of  the  bleach  solution 
contains  0*25  gram  of  the  compound  (10  grams  to  the  litre),  and  the  weight  of 
potassium  iodide  in  10  c.c.  of  the  KI  solution  is  0*5  gram. 

f  Acetic  acid  does  not  liberate  chlorine  from  any  chlorate  which  the  bleach- 
ing powder  may  contain. 

\  Although  this  method  does  not  involve  the  use  of  sodium  thiosulphate,  it 
is  most  conveniently  introduced  at  this  point. 


Valuation  of  Bleaching  Powder.  355 

in  bleaching  powder  by  measuring  the  amount  of  iodine  which  that 
chlorine  is  capable  of  liberating,  it  maybe  determined  by  measuring 
the  amount  of  arsenious  oxide  which  the  oxygen  present  in  the 
active  constituent  of  the  powder  is  capable  of  oxidising.  This 
active  ingredient,  as  stated  above,  is  the  calcium  chloro-hypo- 
chlorite,  Ca(OCl)Cl,  which  contains  I  atom  of  oxygen  to  every  2 
atoms  of  chlorine.  Hence  the  weight  of  bleaching  powder  capable 
of  furnishing  35*5  parts  of  available  chlorine,  will  also  supply  8  parts 
of  oxygen  ;  this  oxygen,  therefore,  may  be  taken  as  the  measure  of 
the  chlorine — 

2Ca(OCl)Cl*  +  As2O3  =  As2O5  +  2CaCl2 

Twenty-five  cubic  centimetres  of  the  freshly  made  and  well- 
shaken  up  bleaching  powder  solution  are  transferred  to  a  small 
beaker  by  means  of  a  pipette,  and  deci-normal  arsenious  oxide 
solution  f  is  gradually  added  from  a  burette.  The  end  of  the  reaction 

*  When  bleaching  powder  is  mixed  with  water,  it  is  converted  into  calcium 
chloride  and  hypochlorite  ;  thus — 

2Ca(OCl)Cl  =  Ca(OCl)2  +  CaCl2 

The  com  position  of  bleaching  powder,  therefore,  maybe,  and  often  is,  expressed 
by  the  formula  Ca(OCl)2,CaCl2,  and  it  will  be  obvious  that  the  proportion 
which  the  oxygen  bears  to  the  chlorine  is  the  same  in  both  cases,  i.e.  8  parts 
by  weight  of  O  to  35^5  parts  of  Cl.  The  compound  calcium  hypochlorite 
Ca(OCl)«j  itself  contains  all  the  oxygen  but  only  half  the  chlorine  originally 
present  in  the  bleaching  powder. 

Whether  or  not  the  decomposition  of  the  bleaching  powder  into  hypochlorite 
and  chloride  by  the  agency  of  water  is  complete,  is  not  known  with  certainty, 
probably  it  is  not.  Neither  is  the  exact  modus  operandi  of  the  water,  in  bring- 
ing about  this  decomposition,  clearly  understood.  Probably  a  cycle  of  changes 
takes  place,  which  may  be  explained  by  the  following  equations ; — 

Ca(OCl)Cl  HC1  ^OCl 

+  H2O  =          +  Ca<         +  CaO 
Ca(OCl)Cl  HC1  XOC1 

HOC1 

=  CaCl2  +  +  CaO 

HOC1 

~~"ocT~ 

=  CaCl2  -f  Ca/         +  H2O 
XOC1 

f  The  deci-normal  solution  used  for  estimating  the  value  of  bleaching 
powder  must  either  be  freshly  prepared,  as  described  on  p.  350,  or  else  it  must 
be  made  with  normal  sodium  carbonate  instead  of  the  bicarbonate,  using  20 
grams  in  place  of  the  30  grams  of  the  latter  salt.  It  has  been  found  in  practice 
that  the  solution  made  from  the  bicarbonate  undergoes  some  change  ort  being 
kept  (the  precise  nature  of  which  is  not  known),  which  introduces  errors  into 
the  analysis  greater  than  are  caused  by  the  action  of  the  normal  carbonate  upon 
the  starch  iodide. 


356  Volumetric  Analysis. 

is  determined  by  means  of  starch  and  potassium  iodide,  used  as  an 
outside  indicator.  After  each  addition  of  the  arsenious  solution, 
the  mixture  is  stirred,  and  a  drop  of  it  removed  upon  a  glass  rod, 
and  brought  into  contact  with  a  piece  of  filter-paper,  which  has  been 
previously  impregnated  with  a  mixture  of  starch  and  potassium 
iodide.*  So  long  as  undecomposed  bleaching  powder  is  present, 
the  liquid  will  cause  a  blue  stain  upon  the  paper.  The  arsenious 
solution  is  cautiously  added,  until  a  drop  of  the  liquid  brought  into 
contact  with  potassium-iodide-and-starch  paper  gives  no  blue 
colour. 

If  the  exact  point  is  overstepped,  a  little  additional  excess  of  the 
arsenious  solution  may  be  added,  together  with  a  few  drops  of  starch, 
and  the  excess  titrated  by  means  of  deci-normal  iodine  solution, 
until  the  blue  colour  is  produced.  On  deducting  this  excess  from 
the  total  volume  of  the  arsenious  solution  used,  the  exact  volume 
which  was  oxidised  by  the  bleaching  powder  is  obtained,  and  from 
this  the  percentage  of  available  chlorine  may  be  calculated. 

(3)  Estimation  of  Manganese  Dioxide.— When  man- 
ganese dioxide  is  heated  with  hydrochloric  acid,  an  amount  of 
chlorine  is  evolved  which  is  the  chemical  equivalent  of  the  avail- 
able oxygen  in  the  manganese  dioxide  (see  p.  337).  The  chlorine 
so  evolved  may  be  absorbed  by  means  of  a  solution  of  potas- 
sium iodide,  whereby  an  amount  of  iodine  equivalent  to  the 
chlorine  (and  therefore  equivalent  to  the  oxygen)  is  liberated.  The 
iodine  thus  set  free  can  then  be  determined  by  means  of  sodium 
thiosulphate. 

About  0*5  gram  of  the  manganese  dioxide,  in  fine  powder,  is 
weighed  out  into  a  small  flask  (Fig.  73),  about  80  to  100  c.c. 
capacity.  The  flask  is  provided  with  a  delivery  tube,  which  is  fitted 
either  by  means  of  a  good  cork  which  has  been  previously  soaked 
in  melted  paraffin  wax,  or  a  rubber  stopper  which  has  been  boiled 
in  caustic  soda  in  order  to  remove  sulphur  ;  or  the  tube  is  made 
by  drawing  out  a  wider  tube,  the  wide  portion  of  which  is  slightly 
conical,  and  has  pushed  over  it  a  short  piece  of  black  rubber  tube. 
This  constitutes  a  perfect  fitting  stopper  and  exit  tube  in  one.  The 
delivery  tube,  bent  as  shown,  passes  through  a  cork  into  a  U-tube 
placed  in  a  beaker  of  cold  water.  The  U-tube  contains  moderately 
strong  potassium  iodide  solution.  (2*5  grams  of  KI  for  0*5  gram 

*  This  paper  is  readily  prepared  by  adding  a  small  quantity  of  potassium 
iodide  to  a  little  clear  starch  solution,  and  dipping  strips  of  unsized  paper 
into  the  mixture.  If  these  strips  are  then  hung  up  to  dry,  they  may  be  pre- 
served indefinitely  in  a  stoppered  bottle. 


Estimations  by  means  of  Iodine  and  Thiosnlphate.  357 

of  MnO2  will  ensure  a  suitable  excess  *)  A  few  fragments  of 
magnesite  are  introduced  into  the  flask  along  with  the  manganese 
ore,  and  about  25  to  30  c.c.  of  strong  hydrochloric  acid  added. 
The  delivery  tube  is  then  attached,  and  a  gentle  heat  applied  to 
the  flask.  The  object  of  the  magnesium  carbonate  is  to  furnish  a 
slow  stream  of  carbon  dioxide  during  the  entire  operation  (magnesite 
being  only  slowly  dissolved  by  the  acid),  which  not  only  sweeps 
out  the  chlorine,  but  prevents  the  liquid  in  the  absorption  tube 
from  being  sucked  back  into  the  flask  during  the  last  stages  of  the 


FIG.  73. 

operation  when  the  acid  is  being  boiled.  As  the  chlorine  enters 
the  potassium  iodide  solution  it  is  wholly  absorbed,  so  long  as  the 
evolution  of  the  gas  is  not  allowed  to  proceed  too  fast.  When  the 
manganese  dioxide  has  all  passed  into  solution,  the  temperature  is 
raised  to  the  boiling-point,  and  the  liquid  allowed  to  boil  for  a  few 
minutes.  When  the  distillation  is  completed,  the  U-tube  is  dis- 
connected, and  the  contents  washed  out  and  diluted  to  100  c.c.  in 

*  A  second  U-tube  is  sometimes  connected  to  the  first  to  serve  as  a  guard, 
lest  any  chlorine  should  escape  unabsorbed  from  the  first.  This,  however,  will 
not  happen  unless  the  distillation  is  made  far  too  rapidly. 


358  Volumetric  Analysis. 

a  graduated  flask.  Twenty-five  cubic  centimetres  of  this  solution 
are  withdrawn  by  means  of  a  pipette,  and  the  liberated  iodine  is 
then  estimated  by  titration  with  deci-normal  sodium  thiosulphate. 
As  soon  as  the  brown  colour  of  the  liquid  fades  to  yellow,  a  little 
starch  is  added,  and  the  addition  of  the  thiosulphate  continued 
until  the  blue  colour  is  just  discharged. 

EXAMPLE. — Weight  of  manganese  dioxide  taken  =  0*495  gram. 
Iodine  solution  diluted  to  100  c.c. ;  25  c.c.  of  the  liquid 

N 
required  of  —  thiosulphate  (i  c.c.  =  0-0127  gram  I)  ...     21  c.c. 

0*0127  x  21  =  0*2667  gram  iodine  in  25  c.c. 
and  0*2667  X  5  =  1*333-5  gram  iodine  liberated  by  0^495  gram 

of  the  ore 
From  the  equation  MnO2  +  4HC1  =  MnCl2  +  2H2O  +  C12  (=  I2), 

127  parts  of  iodine  are  liberated  by  43*5  parts  of  manganese 
dioxide — 

Therefore  127  :  1*3335  • •  43'5  '  °*45674  =  grams  of  MnO2 

and  0-45674  x  IPO  =  92.27  =  percentage  Qf  Mn02    -n 

the  ore 

A  number  of  other  oxygenated  compounds  are  susceptible  of 
being  estimated  by  the  same  method,  namely,  by  distillation  with 
hydrochloric  acid,  and  causing  the  evolved  chlorine  to  liberate  its 
equivalent  of  iodine.  Thus,  a  chromate  treated  in  this  manner 
yields  3  atoms  of  chlorine  for  each  (CrO3)  group  according  to  the 
equations — 

K2CrO,  +  8HC1  =  4H2O  +  2KC1  +  (CrCl8)  +  3C1 
KaCraO7  +  HHC1  =  7H2O  +  2KC1  +  2(CrCl3)  +  6C1 

Similarly,  a  chlorate,  when  treated  with  hydrochloric  acid,  yields 
a  mixture  of  chlorine  and  chlorine  peroxide,  which  is  capable  of 
liberating  iodine  to  the  extent  of  6  atoms  of  iodine  for  each  (C1O3) 
group  ;  thus — 

4KC103  +  I2HC1  =  4KC1  +  6H2O  +  gCl  +  3C1O2 
9C1  +  3C1O2  +  24X1  =  I2KC1  +  6K2O  +  24!^ 

Or  the  relation  of  the  chlorate  to  the  available  oxygen  may  be  more 
clearly  seen  by  the  following  equations  : — 

(1)  KC1O3  =  KC1  +  30 

(2)  30  +  6HC1  =  3H2O  +  6C1 

In  the  case  of  bromates  and  iodates,  only  4  atoms  of  chlorine 
are  disengaged  for  each  group  (Br03)  and  (IO3),  and  therefore  only 


Estimations  by  means  of  Iodine  and  Thiosulphate.    359 

4  atoms  of  iodine  are  liberated  when  the  evolved  gas  is  passed  into 
potassium  iodide.  The  reason  for  this  difference  is  due  to  the  fact 
that  one-third  of  the  chlorine  is  taken  up  in  decomposing  the  mole- 
cule of  potassium  bromide  (or  iodide),  forming  potassium  chloride 
and  the  monochloride  of  the  halogen,  which  remain  behind  in  the 
distilling-flask  ;  thus  — 


(i) 

(2)  30  +6HC1  =  6C1  +  3H2O 

(3)  KI  +  6CI  =  KC1  +  IC1  +  4C1 


SECTION    IV. 
VOLUMETRIC  METHODS  BASED  UPON  PRECIPITATION. 

THE  number  of  volumetric  processes  which  fall  into  this  section 
is  extremely  large.  In  the  following  typical  examples  some  depend 
upon  precipitation  alone,  while  others  are  based  upon  precipitation 
in  conjunction  with  one  of  the  other  methods  described  in  the  fore- 
going sections. 

I.  PRECIPITATION  BY  MEANS  OF  SILVER  NITRATE. 

Deci-normal  Silver  Nitrate  Solution. 

( 1 6*966  grams  o/AgNO3far  litre.") 

4-2415  grams*  of  pure  silver  nitrate  are  exactly  weighed  out 
into  a  250-c.c.  flask,  and  dissolved  in  water.  The  solution  is  made 
up  to  the  quarter  litre  with  cold  water. 

One  cubic  centimetre  of  this  solution  contains  0*010766  gram 
Ag,  and  is  equivalent  to  0*00355  gram  Cl. 

(i)  Estimation  of  Chlorine  in  Sodinm  Chloride.— 
About  i  gram  of  pure  sodium  chloride  is  weighed  out  and  dis- 
solved in  water,  the  solution  being  diluted  up  to  250  c.c. 

Fifty  cubic  centimetres  of  this  solution  are  transferred  to  a 
beaker  by  means  of  a  pipette,  and  three  or  four  drops  of  a  solution 
of  potassium  chromate  f  added  to  serve  as  indicator.  The  deci- 
normal  silver  solution  is  then  gradually  run  in  from  a  burette. 

As  soon  as  the  whole  of  the  chlorine  is  precipitated,  in  the  form 
of  silver  chloride,  any  additional  silver  nitrate  interacts  with  the 
chromate,  precipitating  red  silver  chromate.  As  soon,  therefore, 
as  a  drop  of  the  silver  nitrate  produces  a  permanent  reddish  tinge 
in  the  liquid,  the  titration  is  complete. 

*  Unless  a  number  of  determinations  are  to  be  made,  it  is  not  necessary  to 
prepare  more  than  a  quarter  of  a  litre  of  the  silver  solution. 

f  The  neutral  or  yellow  chromate  K2CrO4.  For  this  titration  the  solutions 
must  be  neutral  and  cold 


Volumetric  Precipitation  Methods.  361 

The  process  may  be  once  or  twice  repeated  for  the  sake  of 
practice ;  and  in  order  to  gain  experience  in  detecting  the  first 
appearance  of  the  red  colour,  a  second  beaker,  containing  about 
an  equal  volume  of  sodium  chloride  and  the  chromate  indicator, 
to  which  a  few  drops  of  silver  nitrate  solution  have  been  added, 
may  be  placed  alongside  the  one  in  which  the  titration  is  being 
made,  so  as  to  compare  the  colours.* 

Alternative  Method. — The  precipitation  of  chlorine  by  silver 
nitrate  (or  vice  versa,  of  silver  by  means  of  a  soluble  chloride)  is 
one  of  the  very  few  examples  of  a  precipitation  process  in  which 
the  end  reaction  can  be  determined  without  recourse  to  a  third 
substance  as  indicator.  This  is  due  to  the  peculiarities  of  the 
precipitate,  which  allow  of  its  ready  subsidence,  leaving  a  clear 
liquid  after  being  shaken.  Moreover,  the  chromate  indicator  is 
inadmissible  in  the  presence  of  acids  (silver  chromate  being  soluble) ; 
hence,  before  it  can  be  employed,  the  solution  to  be  titrated,  if  acid, 
must  be  first  carefully  neutralised  with  sodium  carbonate.  The 
following  method  is  carried  out  in  an  acid  solution  : — 

Fifty  cubic  centimetres  of  the  solution  of  sodium  chloride  pre- 
pared as  above,  are  transferred  to  a  well-fitted  stoppered  bottle 
(narrow  mouth),  and  a  few  drops  of  strong  nitric  acid  added.  The 
bottle  should  be  screened  from  bright  daylight  during  the  process. 
This  is  conveniently  done  by  rolling  a  strip  of  black  paper  (or 
common  brown  paper)  into  a  short  cylinder  about  the  same  height 
as  the  bottle,  and  into  which  the  bottle  will  just  fit ;  or  the  bottle 
may  be  wrapped  in  a  black  cloth.  The  deci-normal  silver  is  then 
added  from  a  burette,  a  considerable  proportion  of  the  total 
quantity  required  being  run  in  at  once.  The  bottle  is  then  briskly 
shaken  without  exposure  to  daylight,  and  the  precipitate  allowed  to 
settle.  If  any  floats  or  adheres  to  the  sides,  it  can  be  made  to  sink 
by  gently  tapping  the  bottle  upon  the  table.  As  soon  as  the  liquid 
has  clarified,  more  of  the  silver  is  added,  and  the  shaking  repeated. 
This  is  repeated  until  the  addition  of  the  silver  solution  causes  no 
further  precipitation.  As  the  point  is  reached  when  the  precipita- 
tion is  nearly  complete,  the  solution  clarifies  less  readily  ;  hence, 
until  experience  is  gained,  the  exact  point  of  completion  is  liable  to 
be  overstepped.  It  is  desirable,  therefore,  when  the  precipitation 

*  Instead  of  the  chromate  indicator  being  employed,  the  chlorine  may  be 
completely  precipitated  by  the  addition  of  an  excess  of  the  deci-normal  silver, 
a  few  drops  of  nitric  acid  being  first  added.  The  mixture  is  then  filtered,  and 
the  excess  of  silver  present  in  the  solution  is  determined  by  titration  with 
ammonium  thiocyanate,  with  a  ferric  salt  as  indicator  (see  p.  365).  Bromides 
and  iodides  may  be  determined  in  the  same  manner. 


362  Volumetric  Analysis. 

is  judged  to  be  complete,  to  add  a  few  drops  of  deci-normal  sodium 
chloride  from  a  burette  (see  p.  364),  which,  if  excess  of  silver  has 
been  added,  will  cause  a  further  turbidity  in  the  liquid.  The 
amount  of  silver  equivalent  to  the  deci-normal  sodium  chloride 
thus  added,  is  deducted  from  that  which  has  been  used. 

In  cases  where  the  chloride  is  present  only  in  very  small 
quantities,  as  in  the  estimation  of  chlorides  in  natural  waters,  the 
analysis  may  be  carried  out  by  the  first  of  the  methods  here  given. 
If  the  second  process  be  employed,  a  carefully  measured  volume 
(about  20  c.c.)  of  deci-normal  sodium  chloride  should  be  added  to 
50  c.c.  of  the  water,  and  the  mixture  titrated  with  the  deci-normal 
silver.  The  volume  of  the  silver  solution  equivalent  to  the  amount 
of  added  chlorine,  deducted  from  the  total  silver  solution  used, 
gives  the  amount  of  silver  required  to  precipitate  the  chlorine 
originally  present  in  the  water. 

(2)  Estimation  of  Cyanogen  in  Potassium  Cyanide.  - 

When  excess  of  silver  nitrate  is  added  to  potassium  cyanide, 
the  reaction  which  takes  place  is  analogous  to  that  between  silver 
nitrate  and  potassium  chloride,  silver  cyanide  being  precipitated — 

KCN  +  AgNO3  =  AgCN  +  KNO3 

N 
And  therefore  I   c.c.  —  silver  nitrate,  containing  0-01766  gram  Ag, 

is  equivalent  to  o'oo2<5  gram  (CN).  Silver  cyanide,  however,  is 
soluble  in  potassium  cyanide,  molecule  for  molecule,  forming  the 
double  cyanide  (soluble)  KCN, AgCN  ;  therefore,  when  silver 
nitrate  is  gradually  added  to  a  solution  of  potassium  cyanide,  the 
first  action  is  the  formation  of  this  double  salt ;  thus — 

2KCN  +  AgNO3  =  KCN,AgCN  +  KNO3 

As  soon  as  the  whole  of  the  potassium  cyanide  has  been  converted 
into  this  soluble  double  salt,  the  addition  of  more  silver  results  in 
its  decomposition  and  the  precipitation  of  silver  cyanide  ;  thus — 

KCN,AgCN  +  AgN03  =  2AgCN  +  KNO3 

Hence  the  completion  of  the  first  stage  can  be  taken  as  the  end 
reaction,  which  will  be  indicated  by  the  first  appearance  of  a 
permanent  precipitate. 

N^ 

For  the  completion  of  the  first  stage,  I  c.c.  —  silver  nitrate  will 

be  equivalent  to  0*0026  x  2  =  0^0052  gram  (CN)  or  0-0130  gram 
KCN. 

The  first  step  in  the  reaction  between  silver  nitrate  and  potas- 
sium cyanide  is  not  interfered  with  by  the  presence  of  soluble 


Precipitations  by  means  of  Silver  Nitrate.      363 

halides  (chlorides,  bromides,  iodides).  If,  therefore,  a  chloride  be 
added  to  the  solution,  no  silver  chloride  is  formed  until  the  cyanide 
has  been  entirely  converted  into  the  double  salt.  When  that  stage 
is  complete,  then  the  further  addition  of  silver  nitrate  results  in  the 
precipitation  of  silver  chloride.  In  fact,  the  completion  of  the  first 
reaction  is  more  distinctly  seen  by  introducing  a  few  drops  of 
sodium  chloride  solution  to  serve  as  indicator,  than  by  depending 
on  the  decomposition  of  the  double  cyanide  and  the  precipitation 
of  silver  cyanide. 

About  0*5  gram  of  potassium  cyanide  is  weighed  out  and  dis- 
solved in  water  in  a  loo-c.c.  flask,  and  the  liquid  diluted  up  to  the 
graduation.  Twenty-five  cubic  centimetres  of  this  are  transferred 
to  a  small  beaker  by  means  of  a  pipette,  and  deci-normal  silver 
nitrate  added  from  a  burette  until  the  first  indications  of  a  faint 
permanent  precipitate  appear.  The  beaker  should  be  placed  upon 
a  dead  black  surface.  A  second  portion  may  be  similarly  titrated 
after  the  addition  of  one  or  two  drops  of  a  dilute  solution  of  sodium 
chloride  to  serve  as  the  indicator. 

(3)  Indirect  Estimation  of  Nitric  Acid  in  Nitre.*— 

If  potassium  nitrate  is  evaporated  with  strong  hydrochloric  acid, 
the  radical  (NO3)  is  exchanged  for  (Cl)  ;  that  is  to  say,  nitric 
acid  is  expelled,  and  potassium  chloride  is  left — 

KNO3  +  HC1  =  KC1  +  HNO3 

The  residue,  after  being  heated  for  a  short  time  to  about  120°,  is 
dissolved  in  water,  the  solution  diluted  to  a  definite  volume,  and 
aliquot  portions  titrated  with  deci-normal  silver,  using  the  chromate 
indicator.  The  chlorine  which  is  estimated  is  thus  the  measure  of 
the  nitric  acid,  35*5  parts  of  Cl  displacing  62  parts  of  (NO3)  or  54 

N 
parts  of  (N2O5).     Therefore  i  c  c.  —  silver  is  equivalent  to  o'oo62 

gram  (NO3)  or  o'oo54  gram  (N2O5). 

The  nitrate  of  any  metal  which  is  thus  capable  of  being  con- 
verted into  a  chloride,  and  whose  chloride  can  be  heated  without 
undergoing  decomposition  or  volatilisation,  may  be  determined  in 
this  way. 

(4)  Indirect  Estimation  of  Calcium  in  Calcium  Car- 
bonate.— The  calcium  carbonate  is  dissolved  in  hydrochloric  acid, 
and  the  solution  evaporated  to  dryness.    The  residue  is  dissolved  in 
water  and  again  evaporated  to  dryness,  and  finally  heated  to  about 
120°  to  expel  the  excess  of  acid.     The  residue  consisting  of  calcium 
chloride  is  dissolved  in  water,  diluted  to  a  definite  volume,  and  the 
chlorine  estimated  in  an  atliquot  portion  by  means  of  silver.     The 

*  The  four  following  examples  serve  to  indicate  the  various  ways  in  which 
indirect  estimations  can  be  made  by  means  of  standard  silver  nitrate. 


364  Volumetric  Analysis. 

chlorine  is  thus  the  measure  of  the  calcium,  35*5  parts  of  chlorine 
being  equivalent  to  20  parts  of  calcium  — 

CaCO3  +  2HC1  =  CaCl2  +  H2O  +  CO, 

N 
Hence  I  c.c.  —  silver  is  equivalent  to  oxx>2O  gram  of  Ca. 

(5)  Indirect  Estimation  of  Carbon  Dioxide  in  Sodium 
Carbonate.  —  The  sodium  carbonate  is  dissolved  in  water,  and  pre- 
cipitated with  barium  chloride.     The  barium  carbonate  is  filtered 
and  thoroughly  washed,  and  then  dissolved  upon  the  filter  in  dilute 
hydrochloric  acid.     The  liquid  is  thoroughly  washed  through  the 
filter  and  evaporated  to  dryness  at  least  twice  (as  directed  above), 
and  afterwards   heated   to   about    120°.     It   is   then  dissolved  in 
water,  and  made  up  to  a  definite  volume.     An  aliquot  portion  is 
taken  and  a  slight  excess  of  sodium  sulphate  added,  in  order  to 
precipitate  the  barium.     The  chromate  indicator  is  then  added  (the 
precipitate  of  barium  sulphate  exerting  no  disturbing  influence  upon 
the  titration),  and  deci-normal  silver  run  in  until  the  red  colour 
appears  — 

Na2CO3  +  Bad.,  =  2NaCl  +  BaCO3 
BaCO3  +  2HC1  =  BaCl2  +  CO2  +  H2O 

The   chlorine   estimated   is   the   measure  of  the  carbon  dioxide, 
2   atoms  of  Cl   being  equivalent  to  I   molecule  of  CO2  ;  or  35  '5 

N 
parts  of  Cl  are  equivalent  to  22  parts  of  CO2.     Hence  I  c.c.  — 

silver  is  equivalent  to  0*0022  gram  of  CO.2. 

(6)  Indirect    Estimation    of  Cadmium   in   Cadmium 
Sulphate.  —  The  sulphate  is  converted  into  carbonate  by  precipi- 
tation with  sodium  carbonate.     The  precipitate  is  filtered,  washed, 
and  dissolved  in  hydrochloric  acid.     The  solution  is  evaporated  to 
dryness  once  or  twice,  and  the  residue  heated  to  expel  the  excess  of 
acid.     It  is  then  dissolved,  and  the  chlorine  in  it  estimated  by  deci- 
normal  silver  with  chromate  indicator  — 


CdSO4  +  Na,CO3  =  CdCO3  +  Na2SO4 
CdCO3  +  2HC1  =  CdCl2  +  CO2  +  H2O 

Chlorine  is  thus  the  measure  of  cadmium,  35*5  parts  of  Cl  being 

N 
equivalent  to  56  parts  of  cadmium.     Therefore  i   c.c.  —  silver  is 

equivalent  to  0*0056  gram  Cd. 

II.  PRECIPITATION  BY  MEANS  OF  SODIUM  CHLORIDE. 
Deci-normal  Sodium  Chloride  Solution. 

(5-85  grams  NaCl  per  litre.} 

5*85  grams  of  purified  sodium  chloride  are  weighed  out  and 
dissolved  in  water,  the  solution  then  being  diluted  up  to  i  litre. 


Ammonium   Thiocyanate.  365 

One  cubic  centimetre  contains  0*00585  gram  NaCl,  or  ©'00355 
gram  Cl,  and  is  equivalent  to  0-010766  gram  Ag. 

Estimation  of  Silver. — The  details  of  the  process  are 
practically  the  same  as  for  the  estimation  of  chlorine  by  means  of 
silver,  p.  360.  When  the  chromate  indicator  is  used  to  determine 
the  end  reaction,  it  should  not  be  added  to  the  silver  nitrate  ;  but 
a  measured  volume  of  the  sodium  chloride  should  be  delivered 
into  a  beaker,  the  indicator  added  to  this,  and  the  silver  solution 
delivered  from  a  burette  until  the  red  colour  makes  its  appearance. 

III.  PRECIPITATION  BY  MEANS  OF  AMMONIUM  THIOCYANATE. 
Deci-uormal  Ammonium  Thiocyanate  Solution. 

(7-6  grams  of  (NHJCNS  per  litre.} 

About  8  grams  of  the  crystallised  salt  (which  is  too  deliquescent 
to  allow  of  exact  weighing)  are  roughly  weighed  out  and  dissolved 
in  a  litre  of  water.  The  solution  is  then  titrated  by  means  of  deci- 
normal  silver  nitrate. 

When  ammonium  thiocyanate  is  added  to  silver  nitrate,  a  white 
precipitate  is  produced,  consisting  of  silver  thiocyanate,  insoluble 
in  nitric  acid — 

AgN03  +  (NH4)CNS  =  Ag-CNS  +  NH4NO3 

This  reaction  takes  place  even  in  the  presence  of  a  ferric  salt, 
and  it  is  not  until  the  whole  of  the  silver  has  been  precipitated  that 
the  characteristic  blood-red  ferric  thiocyanate  makes  its  appearance. 
A  ferric  salt,  therefore  (but  obviously  not  the  chloride),  constitutes 
an  extremely  delicate  indicator  for  this  reaction.  The  most  con- 
venient ferric  salt  is  the  sulphate,  a  solution  of  which  may  be  pre- 
pared by  dissolving  a  few  crystals  of  ferrous  sulphate  in  water  in 
a  boiling-tube,  adding  about  half  the  volume  of  strong  nitric  acid 
(pure),  in  order  to  oxidise  the  iron,  and  then  briskly  boiling  the 
mixture  for  a  few  minutes,  to  expel  all  the  nitrous  fumes.  This 
solution,  which  is  acid  with  excess  of  nitric  acid,  is  then  somewhat 
diluted  with  water,  and  a  measured  volume — say  3  or  4  c.c. — should 
be  uniformly  employed  in  the  titrations.* 

Titration  of  the  Thiocyanate  Solution. — Twenty-five 
cubic  centimetres  of  deci-normal  silver  nitrate  are  transferred  to  a 
small  flask,  and  4  c.c.  of  the  ferric  sulphate  solution  added.f  The 

*  When  a  number  of  titrations  are  to  be  made,  the  ferric  solution  should  be 
made  up  of  some  definite  strength,  and  the  same  volume  of  it  used  for  each 
operation. 

t  The  presence  of  nitric  acid  is  necessary,  and  if  the  ferric  sulphate  solution 
is  prepared  as  described,  it  will  contain  sufficient  acid. 

2    B 


366  Volumetric  Analysis. 

solution  of  the  ammonium  thiocyanate  is  then  run  in  from  a  burette. 
As  each  drop  enters  the  mixture,  the  red  colour  of  ferric  thiocyanate 
appears  for  a  moment,  but  at  once  disappears  on  gently  shaking 
the  flask.  The  solution  is  cautiously  added  until  the  first  indications 
of  a  permanent  red  coloration  are  seen.  From  the  volume  used, 
the  exact  strength  of  the  thiocyanate  solution  is  ascertained,  and 
the  amount  of  dilution  which  it  requires  in  order  to  bring  it  strictly 
to  deci-normal  strength  is  readily  calculated.  After  the  solution 
has  been  thus  adjusted,  it  should  be  once  more  titrated. 

N 
I  c.  c.  —  thiocyanate  contains  0*0076   gram  (NHJCNS,  and 

is  equivalent  to  0*010766  gram  Ag,  0*00355  gram  Cl,  or  0*0127 
gram  I. 

(1)  Estimation  of  Silver. — The  estimation  of  silver  is  con- 
ducted exactly  as   in  the  titration  of  the   standard   thiocyanate 
solution.     The  presence  of  other  metals,  as  lead,  copper,  zinc, 
exerts  no  disturbing  effect  so  long  as  the  solution  contains  free 
nitric  acid.     Hence  the  silver  in  such  an  alloy  as  silver  coinage, 
may  be  determined  directly  in  the  nitric  acid  solution,  by  diluting 
the  liquid  to  a  definite  volume,  and  titrating  an  aliquot  portion  with 
deci-normal  thiocyanate,  and  the  ferric  sulphate  indicator. 

(2)  Indirect    Estimation    of    Chlorine    (Bromine    or 
Iodine). — The  solution  of  the  chloride  is  acidified  with  nitric  acid, 
and  precipitated  by  the  addition  of  a  measured  excess  of  deci-normal 
silver  nitrate.     The  mixture  is  briskly  shaken   and  filtered,  and 
the  precipitate  washed  free  from  silver  nitrate.     The  filtrate  and 
washings  are  then  titrated  with  deci-normal  thiocyanate,  after  the 
addition  of  the  ferric  indicator.     The  excess  of  silver  nitrate  thus 
ascertained,  deducted  from  the  total  volume  taken,  gives  the  amount 
which  was  used  in  precipitating  the  chlorine  ;  and  from  this  the 
percentage  of  the  halogen  is  calculated. 

The  necessity  for  filtering  the  solution  lies  in  the  fact  that  silver 
chloride  is  dissolved  by  alkaline  thiocyanate.  Silver  bromide  is 
less  easily  affected,  while  upon  silver  iodide  it  exerts  practically  no 
solvent  action. 

IV.  PRECIPITATION  BY  MEANS  OF  URANIUM  ACETATE. 

When  a  solution  of  uranium  acetate  or  nitrate  *  is  added  to  a 
solution  of  an  orthophosphate,  in  the  presence  of  sodium  acetate 

*  These  compounds  are  often  spoken  of  as  uranyl  salts,  and  are  represented 
as  containing  the  divalent  group  (UO2)",  analogous  to  the  monovalent  groups, 
bismuthyl,  BiO  (in  basic  bismuth  nitrate,  or  bismuthyl  nitrate,  (BiO)NO3,H2O), 
and  antimony],  SbO  (in  tartar  emetic,  (SbO)K(C4H4O6). 


Precipitation  Methods.  367 

and  acetic  acid,  a  yellow  precipitate  of  uranium  phosphate  is  pro- 
duced, according  to  the  equation — 

(U02)"(C2H302)2  +  HNa2P04  =  2Na(C2H3O2)  +  H(UO2)"PO4 

The  completion  of  the  reaction  may  be  determined  by  means  of 
a  solution  of  potassium  ferrocyanide,  used  as  an  outside  indicator. 
As  soon  as  the  uranium  solution  is  in  excess,  the  ferrocyanide  gives 
a  red-brown  coloration  with  a  drop  of  the  mixture  brought  in 
contact  with  it  upon  a  white  plate. 

Standard  Uranium  Acetate,  (UO2)"(C2H3O2),2H2O.— This 
solution  is  made  of  such  a  strength  that  I  c.c.  is  equivalent  to 
0-005  gram  P2O5. 

From  the  equation  given  above,  it  will  be  seen  that  425-8  parts 
of  the  crystallised  salt  are  equivalent  to  ^  =  71  parts  of  P2O5  ; 
hence— 

71  :  5  : :  425-8  :  29-986  =  weight  of  uranium  acetate  required 
per  litre  * 

About  31  or  32  grams,  therefore,  of  the  crystallised  salt  are 
weighed  out  and  dissolved  in  water  in  a  litre  flask,  with  the 
addition  of  30  to  40  c.c.  of  glacial  acetic  acid.  The  solution  is  then 
diluted  up  to  1000  c.c.  with  water. 

Titration  of  the  Uranium  Solution,  (a)  By  means  of 
HydrogenDisodium  Phosphate. — 10*085  grams  of  pure  recrystallised 
hydrogen  disodium  phosphate  (which  has  not  become  effloresced 
by  exposure  to  the  air)  are  exactly  weighed  out  and  dissolved  in 
water,  the  solution  then  being  made  up  to  one  litre. 

One  cubic  centimetre  of  this  solution  will  contain  0*002  gram, 
P2O5.  It  is  therefore  two-fifths  of  the  strength  of  the  standard 
uranium  solution,  20  c.c.  of  which  should  be  equivalent  to  50  c.c. 
of  the  sodium  phosphate. 

Fifty  cubic  centimetres  of  the  sodium  phosphate  solution  (con- 
taining o- 1  gram,  P2O5)  are  transferred  to  a  beaker  by  means  of  a 
pipette,  and  heated  until  it  just  begins  to  boil  t  (the  titration  is 
best  made  at  a  temperature  about  95°).  The  uranium  solution 
is  then  run  in  from  a  burette,  until,  on  placing  a  drop  of  the  hot 

*  If  uranium  nitrate,  (UO2)"(^O3)2l6H2O,  be  used  instead  of  the  acetate, 
the  exact  weight  of  the  salt  per  litre  to  give  a  solution  of  the  same  value  will 
be  35-47  grams,  the  formula- weight  of  this  salt  being  503-8.  About  36  or  37 
grams,  therefore,  should  be  weighed  out. 

f  If  uranium  nitrate  has  been  employed,  5  c.c.  of  a  lo-per-cent.  solution  of 
sodium  acetate  must  be  added.  This  is  not  necessary  in  tlie  case  of  uranium 
acetate,  as  by  double  decomposition  alkaline  acetate  is  produced  during  the 
titration. 


368  Volumetric  Analysis. 

mixture  upon  a  plate,  and  adding  a  drop  of  potassium  ferrocyanide 
to  it,  a  red-brown  coloration  is  seen.* 

Duplicate  titrations  giving  concordant  results  having  been 
made,  the  exact  value  of  the  uranium  solution  is  ascertained,  and 
if  necessary,  it  may  be  diluted  by  the  addition  of  the  requisite 
amount  of  water  to  bring  it  to  the  correct  strength.  Thus  if  19  c.c. 
instead  of  20  c.c.  were  required  for  the  titration  of  50  c.c.  of  the 
phosphate,  then  every  190  c.c.  must  be  made  up  to  200,  and  the 
solution  titrated  once  more.  If  the  solution  is  more  nearly  exact, 
say  19*5  c.c.  being  required  instead  of  20  c.c.,  it  may  be  employed 
in  this  state,  and  the  volume  used  in  an  analysis  must  then  be 

multiplied  by  the  necessary  factor  ;  in  this  case  — —  =  1-02564. 

(£)  By  means  of  Tricaldum  Phosphate. — In  order  to  accurately 
determine  phosphoric  acid  in  its  combinations  with  lime  or  magnesia 
(as  in  bone-ash,  for  example),  it  is  necessary  that  the  uranium 
solution  should  be  titrated  upon  a  solution  of  calcium  phosphate, 
so  that  the  conditions  of  the  titration  of  the  standard  solution  may 
be  the  same  as  those  obtaining  in  the  actual  analysis. 

The  solution  of  tricalcium  phosphate,  Ca3(PO4)2,  employed  is 
prepared  so  as  to  contain  5  grams  of  the  phosphate  per  litre.  This 
corresponds  very  nearly  to  the  strength  of  the  sodium  phosphate 
solution  above,  and  contains  0*00229  gram  P2O6,  per  cubic  centimetre. 

A  quantity  of  the  purest  commercial  tricalcium  phosphate  is 
weighed  out,  which  contains  exactly  5  grams  of  Ca3(PO4)2.f  This 

*  The  formation  of  the  red-brown  coloration  (uranium  ferrocyanide)  is 
influenced  by  the  presence  of  varying  quantities  of  other  salts  (alkaline  acetates, 
etc.),  which  may  be  present.  It  is  therefore  necessary  to  make  the  conditions 
as  nearly  as  possible  the  same  in  all  cases.  The  same  volumes  of  solutions  of 
approximately  the  same  strength  and  at  the  same  temperature  being  used  ;  and 
the  same  interval  of  time  allowed  for  the  development  of  the  brown  colour 
with  the  indicator.  Again,  if  the  uranium  solution  is  to  be  employed  for 
estimating  phosphoric  acid  in  solutions  which  contain  ammonia  (such  as  in 
urine,  for  example),  a  definite  quantity  of  ammonium  acetate  (5  c.c.  of  a  10- 
per-cent.  solution  to  50  c.c.  of  the  liquid)  should  be  added  to  the  solution  of 
sodium  phosphate  when  standardising  the  uranium  solution. 

f  Even  the  purest  commercial  product  cannot  be  trusted  to  consist  wholly  of 
tricalcium  phosphate,  therefore  it  is  necessary  to  estimate  the  phosphoric  acid 
in  the  particular  sample  used,  by  gravimetric  methods.  For  this  purpose  a 
weighed  quantity,  about  0*4  to  o'6  gram,  is  dissolved  in  nitric  acid,  and  completely 
precipitated  by  means  of  ammonium  molybdate.  After  standing  in  a  warm 
place  some  hours  (conveniently  all  night),  the  precipitate  is  washed  first  by 
decantation,  and  finally  upon  a  filter  with  water  containing  ammonium 
molybdate  and  nitric  acid.  It  is  then  dissolved  in  ammonia,  and  the  phosphoric 
acid  precipitated  as  ammonium  magnesium  phosphate  by  the  addition  of 
"magnesia  mixture,"  and  finally  weighed  as  magnesium  pyrophosphate  (see 
p.  264).  From  the  phosphoric  acid  thus  found  the  percentage  of  Ca3(PO4)2 
in  the  commercial  calcium  phosphate  is  calculated. 


Precipitation  Methods.  369 

is  dissolved  in  the  smallest  quantity  of  hydrochloric  acid,  and  the 
solution  diluted  to  1000  c.c. 

Fifty  cubic  centimetres  of  this  solution  are  then  transferred  to 
a  beaker,  5  c.c.  of  a  lo-per-cent.  solution  of  sodium  acetate,  together 
with  3  or  4  drops  of  acetic  acid,  are  added,  and  the  uranium 
solution  run  in  from  a  burette  until  nearly  the  whole  amount  neces- 
sary has  been  added.*  The  mixture  is  then  heated  to  about  95°,t  and 
the  titration  continued  by  adding  the  uranium  solution  cautiously, 
until  a  drop  of  the  solution  withdrawn  on  a  rod  and  tested  with  potas- 
sium ferrocyanide  gives  the  first  indications  of  red-brown  coloration. 

Estimation  of  Phosphoric  Acid  in  Bone-ash. — About 
1*5  grams  of  bone-ash,  which  has  been  dried  in  a  steam-oven,  are 
weighed  out  and  dissolved  in  the  least  possible  quantity  of  strong 
hydrochloric  acid.  The  solution  is  then  diluted  to  250  c.c.  with 
water.  Fifty  cubic  centimetres  of  this  solution  are  transferred  to 
a  beaker,  and  5  c.c.  of  sodium  acetate  (lo-per-cent.  solution)  added, 
together  with  3  or  4  drops  of  acetic  acid. 

Uranium  acetate  (standardised  by  means  of  calcium  phosphate) 
is  then  added  from  a  burette,  until  the  first  indications  of  colour 
are  perceived  when  a  drop  of  the  mixture  is  tested  on  a  white  plate 
with  a  drop  of  potassium  ferrocyanide.  The  solution  is  then 
heated  to  9o°-95°,  and  again  tested.  If  the  coloration  again 
appears,  the  titration  is  complete  ;  if  not,  the  uranium  solution  is 
added  drop  by  drop,  until  the  liquid  gives  a  faint  colour  with  the 
indicator.  A  second  titration  is  then  made,  in  which  nearly  the 
whole  volume  of  uranium  (which  the  first  experiment  has  shown 
to  be  required)  is  added  at  once  before  the  mixture  is  heated,  and 
the  process  finally  completed  in  the  hot  solution. 

V.  PRECIPITATION  BY  MEANS  OF  SODIUM  SULPHIDE. 

When  a  solution  of  sodium  sulphide  is  added  to  an  ammoniacal 
solution  of  a  zinc  salt,  zinc  sulphide  is  precipitated,  the  end  of  the 

*  When  standardising  the  uranium  solution,  the  volume  required  to  be 
added  is  known  beforehand  within  a  very  near  limit  ;  but  when  estimating 
phosphoric  acid  in  a  phosphate  of  unknown  composition,  the  first  titration  will 
be  made  with  a  view  to  ascertain  approximately  the  volume  of  uranium 
required,  after  which  a  second  experiment  is  made  more  exactly. 

t  It  is  necessary  that  the  solution  should  not  be  heated  until  nearly  the 
whole  of  the  phosphoric  acid  has  been  precipitated  by  the  uranium  solution, 
otherwise  the  calcium  phosphate  itself  will  be  precipitated,  and  will  then  no 
longer  interact  with  the  uranium.  This  difficulty  is  sometimes  evaded  by 
conducting  the  titration  in  the  reverse  way,  i.e.  by  heating  a  measured  volume 
of  the  uranium  solution  in  a  beaker,  and  adding  the  solution  of  the  phosphate 
from  a  burette  until  a  drop  of  the  liquid  just  ceases  to  give  a  coloration  with 
ferrocyanide. 


37O  Volumetric  Analysis. 

reaction  being  indicated  by  means  of  either  an  alkaline  lead 
solution  or  sodium  nitroprusside,  used  as  outside  indicators. 

Standard  Sodium.  Sulphide  Solution. — A  solution  of  con- 
venient strength  should  contain  about  12  grams  of  Na2S  per 
litre,  and  may  be  prepared  as  follows  :  12  grams  *  of  sodium 
hydroxide  are  weighed  out,  and  dissolved  in  50  c.c.  of  water.  One- 
half  of  this  solution  is  then  saturated  with  sulphuretted  hydrogen, 
after  which  the  other  half  is  added,  and  the  mixture  thoroughly 
shaken.  If  the  solution  still  smells  of  the  gas,  a  small  additional 
quantity  of  sodium  hydroxide  is  added,  and  the  liquid  diluted  to 
one  litre, 

Titration  of  the  Sodium  Sulphide  Solution. — The  solu- 
tion of  the  sodium  sulphide  is  titrated  by  means  of  a  standard 
solution  of  zinc  sulphate,  which  is  made  of  such  a  strength  that  I  cc. 
shall  contain  o'oi  gram  Zn.  Such  a  solution  is  prepared  by  weigh- 
ing out  44*15  grams  of  pure  crystallised  zinc  sulphate,  ZnSO4J7H2O, 
dissolving  in  water,  and  diluting  up  to  I  litre.t 

Twenty-five  cubic  centimetres  of  this  solution  are  transferred  to 
a  beaker  by  means  of  a  pipette,  a  small  quantity  of  ammonium 
carbonate  added,  and  then  ammonia,  until  the  precipitate  is  entirely 
redissolved.  The  sodium  sulphide  solution  is  gradually  delivered 
into  this  mixture  from  a  burette,  with  constant  stirring.  The 
progress  of  the  operation  is  watched  by  withdrawing  a  drop  of  the 
liquid  upon  the  end  of  a  fine  glass  rod,  and  bringing  it  into  contact 
with  a  drop  of  sodium  nitroprusside  upon  a  white  plate.  As  soon 
as  the  slightest  excess  of  sodium  sulphide  is  present,  a  violet 
colour  is  imparted  to  the  nitroprusside  solution.  The  alkaline  lead 
solution,  which  may  be  used  as  indicator  in  place  of  the  nitro- 
prusside, is  made  by  adding  a  solution  of  lead  acetate  to  a  mixture 
of  tartaric  acid  with  an  excess  of  caustic  soda,  and  heating  the 
solution  until  it  is  clear.  Drops  of  this  solution  are  placed  at 
intervals  upon  a  piece  of  filter-paper,  and  a  drop  of  the  solution 
undergoing  titration  is  placed  near  to,  but  not  immediately  upon, 
one  of  these.J  As  the  latter  drop  spreads  and  touches  the  lead 

*  The  formula-weight  of  Na2S  and  2NaHO  being  almost  identical,  namely, 
78  and  80. 

f  This  solution  is  a  little  over  three  times  the  strength  of  a  deci-normal 
solution,  which  would  contain  14*35  grams  of  the  salt  per  litre.  An  alternative 
method  of  preparing  the  solution  is  to  dissolve  10  grams  of  pure  zinc  in  the 
smallest  excess  of  dilute  sulphuric  acid,  and  then  make  the  solution  up  to  i 
litre  with  water. 

J  The  reason  for  not  bringing  the  drop  immediately  upon  the  indicator  is, 
that  any  of  the  precipitate  of  ZnS  which  is  brought  out  upon  the  drop  of  the 
mixture  would  be  decomposed  by  the  lead  salt,  with  formation  of  lead  sulphide, 


Precipitation  by  means  of  Sodium  Sulphide.       371 

solution,  a  black  or  brown  stain  of  lead  sulphide  will  be  shown  as 
soon  as  the  sodium  sulphide  is  in  excess.  If  the  sodium  sulphide 
solution  contains  exactly  12  grams  of  Na2S  per  litre;  then  i  c.c. 
(containing  o-oi2  gram  Na2S)  will  be  equivalent  to  0*04415  gram 
of  ZnSO4,7H2O,  or  to  o'oi  gram  of  Zn,  as  deduced  from  the 
following  equation  : — 

ZnSO4,7H2O  +  Na2S  =  ZnS  +  Na2SO4  +  ;H2O 
287  78 

287  :  78  ::  0-04415 : 0-012 
or  65 :  78  : :  0*01 : 0-012 

Hence  25  c.c.  of  the  zinc  solution  would  be  exactly  precipitated  by 
25  c.c.  of  the  sodium  sulphide. 

Estimation  of  Zinc  in  Zinc  Ores   or  Alloys.  —  In  the 

analysis  of  zinc-blende  (p.  279)  or  of  bronze  (p.  274),  the  zinc  may 
be  estimated  by  the  volumetric  process  here  given.  The  solution 
is  rendered  alkaline  by  the  addition  of  excess  of  a  mixture  of 
ammonium  carbonate  and  ammonia,  and  the  standard  sodium 
sulphide  added  from  a  burette,  until  a  drop  of  the  liquid  withdrawn 
upon  a  glass  rod  produces  a  brown  coloration  with  the  alkaline 
lead  indicator  when  applied  in  the  manner  described  above. 

As  the  sodium  sulphide  solution  rapidly  undergoes  change,  it 
should  be  titrated  against  the  standard  zinc  solution  shortly  before 
being  used. 

Other  metals,  e.g.  cadmium  and  copper,  may  also  be  estimated 
by  means  of  sodium  sulphide. 

VI.  CLARK'S  METHOD  FOR  ESTIMATING  THE  HARDNESS  OF 
WATER.* — This  is  a  process  of  precipitation  by  means  of  a  standard 
solution  of  soap.  The  soap  (potassium  oleate)  interacts  with  the 
calcium  and  magnesium  salts  present  in  the  water,  with  the  pre- 
cipitation of  calcium  and  magnesium  oleates.  In  this  process,  the 
soap  solution  itself  serves  as  the  indicator,  not  by  any  colour-change, 
but  by  its  property  of  causing  a  lather  to  form  when  water  contain- 
ing the  solution  is  shaken  up.  So  long  as  the  hardness-producing 
salts  are  undecomposed  by  the  soap,  no  lather  is  formed  on  shaking 
the  mixture,  but  directly  the  soap  solution  is  in  excess  a  permanent 
lather  is  instantly  raised  on  shaking.  By  this  method  of  estimating 
hardness,  the  total  hardness  is  first  determined.  Another  portion 
is  then  boiled  until  the  temporary  hardness  is  destroyed  by  the 

and  the  titration  would  appear  to  be  complete  when  in  reality  it  was  not.     By 
depositing  the  drop  at  some  little  distance  from  the  area  wetted  with  the  lead 
solution,  the  zinc  sulphide  precipitate  is  retained  where  the  drop  first  touches, 
and  only  the  solution  is  able  to  extend,  by  capillarity,  along  the  paper. 
*  For  Hehner's  acidimetric  method,  see  p.  328. 


3/2  Volumetric  Analysis. 

decomposition  of  the  soluble  bicarbonates,  and  the  permanent 
hardness  estimated  in  the  sample.  The  difference  between  the 
two  estimations  represents  the  temporary  hardness. 

Standard  Potassium  Oleate  Solution. — A  strong  solution 
of  this  soap  is  first  prepared  by  rubbing  together  in  a  mortar  75 
grams  of  the  so-called  "  lead  plaster "  of  the  druggist  (plumbi 
emplast.,  Brit.  Pharm.,  which  consists  practically  of  lead  oleate) 
and  20  grams  of  dry  potassium  carbonate.  When  the  two  are 
thoroughly  incorporated,  a  small  quantity  of  methylated  alcohol  is 
added,  and  the  mixture  worked  up  into  the  consistency  of  a  thin 
smooth  cream.  More  spirit  is  then  added,  and  the  contents  of  the 
mortar  rinsed,  by  means  of  spirit,  into  a  bottle,  and  the  mixture 
placed  aside  to  settle.  The  clear  liquid  is  decanted  off  through  a 
filter,  and  the  sediment  finally  washed  upon  the  filter  with  more 
spirit.  The  volume  of  the  liquid  may  now  be  made  up  to  200  or 
250  c.c.  by  adding  a  mixture  of  equal  volumes  of  spirit  and  water. 

Titration  and  Adjustment  of  the  Soap  Solution.— For 
this  purpose  a  standard  solution  of  calcium  chloride  is  prepared, 
containing  0*222  gram  of  CaCl2  per  litre.  This  is  equivalent  to 
0*200  gram  CaCO3  per  litre,  and  50  c.c.  will  therefore  be  equivalent 
to  0*01  gram  of  CaCO3. 

0*2  gram  of  pure  calcite,  or  Iceland  sfiar,  is  exactly  weighed  out 
and  dissolved  in  dilute  hydrochloric  acid  in  a  covered  vessel,  pre- 
ferably a  platinum  dish.  The  solution  is  then  evaporated  to  dryness 
upon  a  steam-bath.  The  residue  is  treated  with  water,  and  the 
solution  again  taken  down  to  dryness.  This  process  should  be  once 
or  twice  repeated  in  order  to  ensure  the  removal  of  all  the  acid. 
The  residue  is  finally  dissolved,  and  the  solution  made  up  to  i  litre. 

Fifty  cubic  centimetres  are  then  transferred  to  a  stoppered 
bottle  of  about  250  c.c.  capacity  by  means  of  a  pipette.*  The 
soap  solution  is  then  added  from  a  burette  I  c.c.  at  a  time, 
and  the  contents  of  the  bottle  briskly  shaken  after  each  addition. 
As  the  point  is  approached  when  a  lather  is  raised,  the  volume 
added  each  time  is  reduced,  and  the  process  continued  until  the 
lather  remains  permanent  for  two  minutes,  the  bottle  being  laid 
upon  its  side. 

From  this  preliminary  experiment,  the  amount  of  dilution  which 
the  soap  solution  will  require,  in  order  that  exactly  14/25  c.c.  shall 
be  necessary  to  produce  the  permanent  lather  with  50  c.c.  of  the 

*  The  contents  of  the  pipette  must  be  allowed  to  flow  out  •without  being 
blown  out  by  the  breath  (see  p.  307),  because  it  is  very  important  to  avoid  intro- 
ducing any  carbon  dioxide  into  the  bottle,  as  this  gas  affects  the  soap  solution. 


Precipitation  Methods.  373 

calcium  chloride  solution,  can  be  calculated.  The  soap  should  then 
be  diluted  by  the  addition  of  a  mixture  of  methylated  spirit  and 
water  in  equal  volumes,  so  as  to  reduce  its  strength  nearly,  but  not 
quite,  to  this  strength— say  to  such  a  strength  that  about  13*5  c.c. 
shall  give  the  lather  with  50  c.c.  of  the  standard  calcium  chloride. 

It  is  then  allowed  to  stand  in  a  stoppered  bottle  for  24  hours, 
during  which  time  a  slight  sediment  usually  settles  out,  and  the 
solution  slightly  loses  strength.  It  is  then  decanted  or  filtered, 
and  titrated  again  against  50  c.c.  of  the  lime  solution,  and  finally 
adjusted  to  the  exact  strength  by  the  addition  of  more  alcohol  and 
water. 

14*25  c.c.  of  the  soap  will  therefore  be  equivalent  to  0*01  gram 
CaCO3  ;  and  o'oi  gram  CaCO3  in  50  c.c.  of  water  equal  20*00  parts 
in  100,000,  or  20  degrees  of  hardness. 

(1)  Estimation  of  the  Total  Hardness  in  Ordinary  Tap- 
water.— Fifty  cubic  centimetres  of  the  water  are  introduced  into  the 
stoppered  bottle  as  above  (see  footnote),  and  briskly  shaken  in  order 
to  expel  dissolved  carbon  dioxide,  and  the  air  in  the  bottle  is  then 
either  swept  out  by  means  of  a  current  of  air  from  a  pair  of  bellows 
(the  blowpipe  bellows),  or  is  sucked  out  by  means  of  a  glass  tube. 
The  soap  solution  is  then  added,   I   c.c.  at  a  time  at  first,  but 
diminishing  to  one  or  two  drops  as  the  operation  proceeds,  with  a 
brisk  shake  up  after  each  addition.     As  soon  as  the  soap  is  in 
excess  and  a  permanent  lather  will  remain  for  two  minutes,  the 
bottle  being  placed  upon  its  side,  the  titration  is  complete,  and  the 
degree  of  hardness  corresponding  to  the  volume  of  soap  solution 
employed  is  ascertained  from  the  table  of  hardness  (see  Appendix). 
If  the  sample  of  water  is  so  hard  that  50  c.c.  require  more  than 
1 6  c.c.  of  the  soap,  the  water  must  be  diluted  by  taking,  say,  half 
this  volume,  and  diluting  it  up  to  50  c.c.  with  distilled  water  which 
has  been  recently  boiled  for  10  minutes  to  expel  carbon  dioxide, 
and  quickly  cooled.     The  degree  of  hardness  which  this  diluted 
sample  contains  is  doubled  in  order  to  express  the  actual  hardness 
of  the  water. 

(2)  Estimation  of  the  Permanent  Hardness  in  Water. 
—  One  hundred  and  fifty  cubic  centimetres  of  the  water  are  placed 
in  a  half-litre  flask  and  counterpoised  upon  a  rough  balance.     The 
water  is  gently  boiled  for  half  an  hour  in  order  to  completely  pre- 
cipitate the  salts  which  cause  the  temporary  hardness.     It  is  then 
cooled  and  replaced  upon  the  balance,  and  cold,  recently  boiled 
distilled  water  added  until  the  original  weight  is  restored.     The 
water  is  then  filtered  (the  filter  being  used  without  being  previously 


374  Volumetric  Analysis. 

wetted),  and  50  c.c.  of  the  clear  filtrate  titrated  with  the  standard 
soap  solution,  as  in  the  previous  case. 

The  difference  between  the  permanent  and  total  hardness  gives 
the  temporary  hardness. 


APPENDIX  TO  SECTION  IV. 

Estimation  of  Copper  by  means  of  Potassium  Cyanide. 

—When  a  solution  of  potassium  cyanide  is  added  to  an  alkaline 
solution  of  a  copper  salt,  a  soluble  double  cyanide  is  formed  (see 
p.  87).  As  this  compound  is  colourless,  while  the  original  copper 
solution  is  blue,  the  end  of  the  reaction  is  readily  seen  by  the  dis- 
appearance of  the  blue  colour. 

Standard  Potassium  Cyanide  Solution.— This  solution 
is  made  of  such  a  strength  that  i  c.c.  shall  be  equivalent  to  0*005 
gram  of  copper.  As  the  solution  cannot  be  made  of  the  exact 
strength  by  weighing  out  the  actual  quantity  of  cyanide  required, 
an  approximate  solution  is  first  prepared  by  dissolving  about  25 
grams  of  potassium  cyanide  in  a  litre  of  water.  The  precise  copper- 
value  of  this  solution  is  then  determined  by  titration  against  a 
standard  copper  solution,  from  which  the  exact  amount  of  dilution 
necessary  to  reduce  it  to  the  required  strength  is  calculated. 

Titration  and  Adjustment  of  the  Cyanide  Solution, 
by  means  of  Standard  Copper  Solution. — Five  grams  of  pure 
electrolytic  copper  are  weighed  out  into  a  litre  flask,  and  dissolved 
in  a  mixture  consisting  of  equal  volumes  of  strong  nitric  acid  and 
water.  The  flask  may  be  gently  warmed  upon  a  steam-bath.  When 
this  metal  is  wholly  dissolved,  the  solution  is  cooled  and  diluted  up 
to  i  litre  with  cold  water. 

One  cubic  centimetre  of  this  solution  will  contain  oxx>5  gram  Cu. 

Twenty-five  cubic  centimetres  of  the  standard  copper  solution 
are  transferred  to  a  beaker,  and  made  alkaline  by  the  addition  of  a 
slight  excess  of  sodium  carbonate  (the  solution  of  sodium  carbonate 
being  cautiously  added,  and  the  beaker  covered,  to  avoid  loss  from 
effervescence). 

To  this  mixture,  in  which  the  precipitate  is  suspended,  is  added, 
by  means  of  a  pipette,  i  c.c.  of  ammonia  solution,  prepared  by 
mixing  I  volume  of  strong  ammonia  (sp.  gr.  o'88)  with  2  volumes 
of  water.  This  dissolves  the  precipitate,  and  produces  a  deep  blue 
solution.*  The  potassium  cyanide  solution  is  then  delivered  into 

*  The  reason  for  neutralising  with  sodium  carbonate  before  adding  ammonia, 
is  to  avoid  the  introduction    of  ammoniacal  salts,   which  by  their  presence 


Estimation  of  Copper  by  means  of  Potassium  Cyanide.  37  5 

this  solution  from  a  burette,  until  the  blue  colour  is  just  discharged. 
The  amount  of  dilution  which  the  cyanide  solution  requires  in  order 
to  reduce  its  strength  to  the  exact  standard,  i.e.  so  that  25  c.c.  of 
the  copper  solution  shall  require  exactly  25  c.c.  of  the  cyanide,  is 
calculated,  and  the  necessary  water  added. 

Thus,  suppose  21  c.c.,  instead  of  25,  were  required  in  the  first 
titration,  then  210  c.c.  must  be  diluted  to  250  c.c.,  or  840  c.c.  will 
be  diluted  to  I  litre. 

Estimation  of  Copper  in  Copper  Ores.3" 

Epitome  of  Process. — The  ore  is  dissolved  in  hydrochloric 
and  nitric  acids,  and  after  filtration  and  dilution,  the  copper  is 
precipitated  by  means  of  a  rod  of  metallic  zinc  in  contact  with  a 
piece  of  platinum.  The  spongy  copper  is  washed,  and  then  dissolved 
in  nitric  acid.  The  solution  is  made  alkaline  with  sodium  carbonate, 
the  precipitate  dissolved  by  the  addition  of  ammonia,  and  the  blue 
liquid  is  finally  titrated  with  standard  potassium  cyanide  in  the  cold. 

Exactly  5  grams  of  the  ore,  in  fine  powder,  are  weighed  into  a 
flask,  and  gently  warmed  with  a  mixture  of  40  c.c.  of  strong  hydro- 
chloric acid  and  10  c.c.  of  water.  Six  cubic  centimetres  of  a  mixture 
of  strong  nitric  acid  and  water  in  equal  volumes  are  then  added, 
and  the  solution  kept  gently  simmering  for  about  half  an  hour,  after 
which  it  is  briskly  boiled  for  about  10  minutes. 

The  insoluble  gangue  is  then  separated  by  filtration,  the  filter 
being  thoroughly  washed,  and  the  copper  precipitated  from  the 
warm  liquid  by  introducing  into  it  a  rod  of  zinc,  upon  which  a 
strip  of  platinum  is  loosely  twisted.  It  is  essential  that  there  shall 
be  an  excess  of  zinc,  the  rod  therefore  should  weigh  about  50  grams, 
and  should  be  as  free  from  lead  as  possible.  In  about  half  an  hour 
the  whole  of  the  copper  will  be  precipitated,  after  which  the  zinc 
rod  is  withdrawn  (leaving  the  platinum  in  the  vessel,  as  a  portion 
of  the  copper  is  deposited  upon  it),  and  rinsed  into  the  beaker.  The 
spongy  copper  is  washed  free  from  zinc  chloride  and  other  dissolved 
salts  by  carefully  decanting  the  liquid,  and  once  or  twice  washing 
the  precipitate  by  decantation.  The  last  wash-water  is  drained 

influence  the  decolorising  of  the  copper  solution  by  the  cyanide.  This  is  some- 
times obviated  to  a  considerable  extent  by  evaporating  the  acid  solution  of  the 
copper  down  to  dryness,  to  expel  the  acid.  In  the  ordinary  process  of  the 
analysis  of  copper  ores,  however,  this  involves  much  loss  of  time,  hence  the  acid 
solutions  are  neutralised  as  here  described,  and  therefore  the  standardisation 
of  the  potassium  cyanide  solution  should  be  conducted  with  a  copper  solution 
under  similar  conditions  as  obtain  in  the  analyses  for  which  the  cyanide  is 
employed. 

*  The  method  here  given  for  the  estimation  of  copper  is  essentially  that 
known  as  the  "  Mansfield,"  or  as  Steinbeck's  (the  discoverer)  process. 


376  Volumetric  Analysis. 

away  as  thoroughly  as  possible,  and  the  spongy  copper  (both  that 
which  is  loose,  as  well  as  that  which  is  deposited  on  the  platinum) 
is  dissolved  by  adding  16  c.c.  of  nitric  acid  (i  volume  strong  acid 
to  i  volume  water).  The  solution  is  then  diluted  to  100  c.c.  with 
water.* 

Fifty  cubic  centimetres  of  this  solution  are  transferred  to  a 
beaker,  made  alkaline  by  the  addition  of  sodium  carbonate,  i  c.c* 
of  ammonia  (i  volume  strong  ammonia  to  2  volumes  of  water) 
added,  and  the  blue  solution  titrated  with  the  standard  cyanide. 
The  number  of  cubic  centimetres  of  the  cyanide  used,  multiplied 
by  2,  gives  the  volume  which  would  be  required  for  the  whole  of 
the  copper  solution  ;  and  since  each  cubic  centimetre  of  the  cyanide 
is  equivalent  to  0*005  gram  Cu,  and  since,  also,  5  grams  of  the 
ore  were  employed  for  analysis,  the  number  of  cubic  centimetres  of 
cyanide,  multiplied  by  0*1,  gives  at  once  the  percentage  of  copper 
in  the -ore.  For  example,  suppose  the  50  c.c.  of  the  copper  solution 
required  41  c.c.  cyanide,  then — 

41  x  2  =  82  cc.  =  cyanide  required  for  the  whole  of 

the  copper  in  the  5  grams  of  ore 
i  c.c.  cyanide  =  0-005  gram  Cu 
therefore  82  x  0-005  —  0*410  gram  Cu  in  5  grams  ore 

and  0>4*Q  x  I0°  =  8-2  per  cent,  of  copper 
or  82  x  o-i  =  8'2 

*  If  the  ore  contains  less  than  6  per  cent,  of  copper,  the  reduced  metal  may 
be  dissolved  in  half  this  volume  of  nitric  acid,  the  solution  diluted  somewhat, 
and  whole  solution  used  for  one  titration. 


SECTION     V. 
GAS  ANALYSIS. 

General  Principles. —The  principles  upon  which  the  various 
methods  of  gas  analysis  are  based  are  themselves  extremely  simple, 
although  in  some  cases  the  apparatus  employed,  and  the  manipu- 
lative details  involved,  are  often  somewhat  complicated  and  intricate. 

This  is  more  especially  the  case  when  the  analysis  is  to  be 
carried  out  with  that  high  degree  of  refinement  and  exactness 
which  is  often  necessary  in  scientific  research.  Until  comparatively 
recent  times  these  exact,  extremely  slow,  and  somewhat  recondite 
methods  of  gas  analysis  were  practically  the  only  processes  avail- 
able ;  but  of  late  years  the  demand  for  simpler  and  more  rapid 
methods,  sufficiently  exact  for  technical  purposes,  has  resulted  in 
the  invention  of  numerous  forms  of  "  gas  apparatus  "  which  are 
capable  of  yielding  good  results  with  the  minimum  expenditure  of 
time  and  trouble.* 

The  methods  by  which  gases  are  estimated  may  be  classified  in 
the  following  order  : — 

(1)  By  Absorption  in  a  Suitable  Reagent,  and  Subse- 
quent Titration. — For  example,  the  carbon  dioxide  present  in 
a  gaseous  mixture,  such  as  the  atmosphere,  may  be  estimated  by 
bringing  a  known  volume  of  the  gas  into  contact  with  an  excess  of 
a  standard  solution  of  barium  hydroxide.      Barium  carbonate  is 
thereby  precipitated,  and  the  excess  of  barium  hydroxide  is  deter- 
mined by  titration  with  standard  oxalic  acid. 

Or,  again,  the  chlorine  present  in  the  exit  gases  from  the 
"  bleach  "  chambers  may  be  estimated  by  aspirating  a  measured 
volume  of  the  gas  through  a  solution  of  potassium  iodide,  and 
subsequently  determining  the  liberated  iodine  by  titration  with 
sodium  thiosulphate. 

(2)  By  Absorption,  and  Subsequent  Measurement  of 

*  For  a  description  of  the  more  refined  and  elaborate  methods  of  gas 
analysis,  the  student  is  referred  to  Button's  "  Volumetric  Analysis."  In  this 
book  only  the  simpler  processes  will  be  described. 


378  Gas  Analysis. 

the  Residual  Gas.— For  example,  the  carbon  dioxide  in  a 
mixture  of  gases  is  determined  by  exposing  a  measured  volume  of 
the  mixture  to  the  action  of  caustic  potash,  and  after  the  whole 
of  the  carbon  dioxide  has  been  absorbed,  the  volume  of  the  residual 
gas  is  measured.  The  difference,  or  the  contraction,  represents 
the  carbon  dioxide  which  was  present. 

(3)  By  the  Combustion  of  the  Gas,  with  the  Sub- 
sequent Measurement  of  the  Contraction,  and  Estima- 
tion of  the  Carbon  Dioxide,  if  any  is  formed. — For 
example,  hydrogen  in  a  gaseous  mixture  may  be  estimated 
by  adding  to  a  known  volume  of  the  gas  a  measured  volume  of 
oxygen  (or  air)  more  than  sufficient  to  combine  with  all  the 
hydrogen.  These  are  caused  to  unite  (by  methods  described  later), 
and  the  "contraction"  ascertained  by  again  measuring  the  gas. 
Since  2  volumes  of  hydrogen  and  I  volume  of  oxygen  unite  to 
form  water  (which  then  practically  occupies  no  volume),  two-thirds 
of  the  contraction  represents  the  hydrogen  originally  present. 

When  the  gas  to  be  estimated  contains  carbon  and  hydrogen 
(as  in  marsh  gas,  ethylene,  etc.),  after  the  contraction  due  to 
combustion  has  been  measured,  the  volume  of  carbon  dioxide 
produced  is  determined  by  absorption  with  caustic  potash  and 
measurement  of  the  residue,  as  in  (2)  above. 

I.  ESTIMATION  OF  GASES  BY  ABSORPTION  AND  SUBSEQUENT 

TlTRATION. 

(a)  Carbon  Dioxide  in  the  Air  (Pettenkofer's  method). 

Epitome  of  Process. — The  sample  of  air  to  be  examined  is 
contained  in  a  large  glass  jar  of  known  capacity,  which  can  be 
closed  with  a  well-fitting  bung.  The  temperature  of  the  gas  and 
the  atmospheric  pressure  are  noted.  A  measured  volume  of  a 
solution  of  barium  hydroxide  of  known  strength  is  introduced  into 
the  jar,  and  thoroughly  shaken  up  with  the  enclosed  air  until  the 
carbon  dioxide  is  entirely  absorbed.  Aliquot  portions  of  the 
solution  are  then  withdrawn,  and  titrated  with  a  standard  solution 
of  oxalic  acid. 

Standard  Oxalic  Acid. — In  order  to  simplify  the  subsequent 
calculations,  the  strength  of  the  oxalic  acid  solution  may  be 
arranged  so  that  I  c.c.  is  equivalent  to  I  c.c.  of  carbon  dioxide 
measured  under  standard  conditions — that  is  to  say,  i  c.c.  of  the 
acid  is  capable  of  saturating  a  quantity  of  barium  hydroxide  which 
would  be  decomposed  by  this  volume  of  carbon  dioxide,  the  weight 
of  v/hich  is  0-00197  gram. 


Estimation  of  Carbon  Dioxide  in  the  Air.      379 

Oxalic  acid  solution  of  this  strength  would  contain  5 '6442  grams 
of  the  crystallised  acid,  H2C2O4,2H2O,  in  one  litre ;  but  since 
dilute  solutions  of  oxalic  acid  are  unstable,  it  is  better  to  prepare 
the  solution  of  ten  times  this  strength,  and  dilute  it  when  required. 

56*442  grams,  therefore,  of  the  pure  dry  crystallised  compound 
are  exactly  weighed  out,  dissolved  in  cold  air-free  water,*  and  the 
solution  made  up  to  one  litre.  For  use,  10  c.c.  are  transferred  to 
a  loo-c.c.  flask  by  means  of  a  pipette,  and  the  solution  diluted  up 
to  loo  c.c.  with  air-free  water. 

Barium  Hydroxide  Solution  (baryta-water). — Forty  to  fifty 
grams  of  crystallised  barium  hydroxide,  Ba(HO)2,8H2O,  are 
roughly  powdered,  and  placed  in  a  large  stoppered  bottle,  and 
a  litre  of  water  added.  The  mixture  is  thoroughly  shaken  from 
time  to  time  until  the  water  is  saturated,  after  which  it  is  allowed 
to  settle.  The  clear  liquid  is  decanted  off  or  filtered  into  a  stop- 
pered bottle,  and  diluted  with  an  equal  volume  of  air-free  water.t 
The  stopper  of  the  bottle  should  be  greased  with  resin  cerate,  so 
as  to  exclude  the  air  as  much  as  possible. 

A  large  glass  bottle  or  jar  $  (whose  mouth  is  sufficiently  wide  to 
admit  the  hand  and  arm,  so  that  it  may  be  readily  wiped  dry  with 
a  glass-cloth)  is  fitted  with  a  good  bung,  preferably  of  rubber.  A 
hole  is  bored  in  the  bung,  and  this  hole  is  closed  with  a  small  cork 
(see  Fig.  74).  The  capacity  of  the  jar  is  ascertained  by  filling  it 
with  water  up  to  the  bung,  and  then  measuring  the  volume  of  the 
water.  The  capacity,  which  should  not  be  less  than  6  litres,  and 
may  with  advantage  be  8  or  10  litres,  should  be  scratched  upon 
the  vessel. 

The  jar  is  filled  with  the  air  to  be  tested,  by  leading  into  it,  right 
to  the  bottom,  a  piece  of  rubber  tube,  which  is  attached  to  a  pair 
of  bellows  (either  hand  or  foot-bellows),  and  blowing  a  stream  of 
air  for  several  minutes  in  order  to  ensure  the  complete  displacement 

*  Distilled  water  which  has  been  recently  boiled  to  expel  carbon  dioxide, 
and  quickly  cooled. 

f  This  solution  will  roughly  approximate  the  same  relative  strength  as  the 
oxalic  acid  above  mentioned.  A  solution  of  exactly  equivalent  strength,  of 
which  i  c.c.  -o-coigygram  CO2  (i.e.  i  c.c.  CO2  at  N.T.P.),  would  contain 
14-11  grams  of  Ba(HO)2,8H2O  per  litre.  Such  a  solution  cannot  be  obtained 
by  direct  weighing,  as,  the  salt  being  efflorescent,  its  state  of  hydration  is 
uncertain ;  and  it  also  absorbs  atmospheric  carbon  dioxide.  A  saturated 
solution  at  the  ordinary  temperature  contains  about  32  grams  per  litre  ;  hence 
if  this  be  diluted  with  its  own  volume  of  water,  the  solution,  as  first  prepared, 
will  contain  about  16  grams  per  litre.  As  the  solution  constantly  undergoes 
change  by  the  absorption  of  atmospheric  carbon  dioxide,  its  value  must  be 
determined  every  time  it  is  used,  by  titration  against  the  standard  oxalic  acid. 

+  The  glass  jars  used  by  confectioners  answer  the  purpose  very  well. 


38o 


Gas  Analysis. 


of  the  air  already  in  the  vessel.  The  bung  is  then  inserted.  The 
temperature  of  the  air  in  the  immediate  vicinity  of  the  bottle  is 
noted,  and  also  the  height  of  the  barometer  at  the  time. 

One  hundred  cubic  centimetres  of  the  baryta-water  are  then 
delivered  into  the  jar  by  means  of  a  pipette,  through  the  hole  in  the 
bung,  and  the  cork  quickly  replaced.  The  liquid  is  then  made  to 
wet  the  interior  surface  of  the  glass  by  slowly  revolving  the  vessel 
upon  its  side  ;  and  it  is  left  in  contact  with  the  gas,  being  shaken 
up  at  intervals,  for  about  half  to  three-quarters  of  an  hour,  by 
which  time  the  whole  of  the  carbon  dioxide  will  be  absorbed. 

While  this  absorption  is  going  on,  the  exact  value  of  the  baryta- 
water  is  determined  by  transferring  25  c.c.  to  a  small  beaker,  and 
adding  the  dilute  standard  oxalic  acid  from  a  burette  until  the 
liquid  ceases  to  give  any  indication  of  a  brown  colour  when  a  drop 
of  it  is  placed  upon  a  piece  of  turmeric  paper  by  means  of  a  glass 
rod.  As  the  alkali  is  gradually  neutralised  by  the  oxalic  acid,  the 
brown  colour,  which  at  first  is  very  evident,  gradually  shows  more 
and  more  as  a  faint  fringe  of  colour  round  the  edge  of  the  moistened 
spot  upon  the  turmeric  paper,  until  finally  it 
disappears. 

Suppose,  for  example,  that  27*5  c.c.  of  the 
oxalic  acid  are  used,  then  100  c.c.  of  this 
baryta  solution  would  require  no'o  c.c.  of  the 
oxalic  acid  to  exactly  neutralise  it. 

When  the  absorption  of  the  carbon  dioxide 
is  complete,  25  c.c.  of  the  turbid  liquid  are 
withdrawn  by  means  of  a  pipette,  to  which  a 
piece  of  narrow  glass  tube  has  been  attached 
in  order  that  it  may  reach  to  the  bottom  of 
the  jar.  This  is  introduced  through  the  hole 
in  the  bung,  in  the  manner  shown  in  the  figure, 
and  when  the  pipette  is  full,  the  piece  of  tube 
is  detached,  and  the  liquid  allowed  to  drip 
from  the  pipette  until  it  reaches  the  graduation 
mark.  The  measured  volume  is  then  trans- 
ferred to  a  beaker,  and  titrated  with  the  oxalic 
acid  as  quickly  as  possible,  so  as  to  expose  it 
to  the  air  for  the  shortest  time.  The  end  of 
the  reaction  is  indicated  by  means  of  turmeric 
FlG'  74'  paper  used  as  described  above.* 

*  The  oxalic  acid  has  no  action  upon  the  precipitated  barium  carbonate. 
Some  chemists  use  a  dilute  standard  hydrochloric  acid  instead  of  oxalic  acid  ; 

either  a  —  solution,  or  one  arranged  so  that  i  c.c.  is  equivalent  to  o'ooi  gram 


Estimations  by  A  bsorption  and  Subsequent  Titration.     3  8 1 

A  duplicate  titration  should  be  made  with  a  second  portion  of 
the  liquid. 

Since  25  c.c.  (out  of  the  100  c.c.  of  baryta-water  originally 
placed  in  the  jar)  have  been  used  for  the  titration,  the  volume  of 
oxalic  acid  used,  multiplied  by  four,  will  give  the  amount  of  oxalic 
acid  required  to  neutralise  the  100  c.c.  of  baryta- water  after  the 
absorption  of  the  carbon  dioxide  0in  the  known  volume  of  air. 
This  obviously  will  be  less  than  that  required  by  100  c.c.  of  the 
original  baryta-water  by  just  the  volume  of  carbon  dioxide  which 
was  contained  in  the  sample  of  air. 

EXAMPLE. — 

Capacity  of  glass  jar  =  10-5  litres 
temperature  of  sample  =  16° 

barometer  reading '=  767  mm.. 

then  ^  *  273  ^767  =  10>01  litreg  =  volume  of  the 
sample  reduced  to  N.T.P. 

Value  of  the  baryta  solution — 

ico  c.c.  =  1 10  c.c.  of  the  standard  oxalic  acid 

100  c.c.  of  the  baryta  used  for  absorption  ;  25  c.c.  taken  out  for 
titration. 

25  c.c.  required  26-45  c  c.  standard  oxalic  acid 
.*.  ico  c.c.  would  require  26-45  x  4  =  105-80  c.c.  oxalic  acid, 
no  —  I05'8o  —  4-2  c.c.  =  the  volume  of  oxalic  acid  which  is  equi- 
valent to  the  carbon  dioxide  absorbed   by  the 
loo  c.c.  of  baryta-water 

But  since  i  c.c.  oxalic  acid  =  i  c.c.  CO2  at  N.T.P. 

4*2  c.c.        „          =  4-2  c.c.  CO2  present  in  lo'oi  litres  (or 
10,010  c.c.)  of  air 

therefore—  -  —  0'0419  =  percentage   of    CO2    by 

(b)  Estimation  of  Sulphur  Dioxide  in  Furnace  Gases. 

Epitome  of  Process.— By  means  of  a  pipe  inserted  into  the 
flue,  a  measured  volume  of  the  furnace  gases  is  aspirated,  with  a 
suitable  aspirator,  through  a  known  volume  of  a  dilute  standard 
solution  of  iodine.  The  excess  of  iodine  is  then  titrated  with  a 
standard  solution  of  sodium  thiosulphate. 

The  standard  solutions  used  for  this  estimation  may  be  the 
—  solutions  described  on  p.  348,  in  which  case  i  c.c.  of  the  iodine 

CO2.  Hydrochloric  acid  of  such  a  strength  that  i  c.c.  shall  be  equivalent  to 
0-00197  gram  CO2  (i.e.  i  c.c.  CO2)  cannot  be  used,  as  this  would  act  upon  the 
precipitated  barium  carbonate. 


382 


Gas  Analysis. 


(and  therefore  indirectly  i  c.c.  of  the  thiosulphate)  is  equivalent  to 
0*0032  gram  of  sulphur  dioxide.  It  is  more  convenient,  however, 
and  simplifies  the  calculations,  to  employ  solutions  of  such  a 
strength  that  i  c.c.  shall  equal  i  c.c.  SO2  measured  at  N.T.P.,  i.e. 
0*002867  gram  SO2  instead  of  0-0032  gram. 

Such  solutions  will  contain  n'379  grams  of  iodine,  and  22*22 
grams  of  sodium  thiosulphate  joer  litre  respectively,  and  they  may 
be  prepared  either  by  weighing  out  these  quantities,  or  by  diluting 

N  N 

the  ordinary  —  solutions  ;  TOO  c.c.  of  the  —  solution   (both  the 

iodine  and  the  thiosulphate)  being  diluted  to  111*6  c.c.* 

In  cases  where  the  percentage  of  sulphur  dioxide  in  the  gas 
under  examination  is  very  small,  it  is  better  to  employ  solutions  of 
one-tenth  this  strength,  in  which  i  c.c.  is  equivalent  to  0*1  c.c.  SO2. 


FIG.  75. 

One  hundred  cubic  centimetres  of  the  standard  iodine  solution 
are  placed  in  a  flask  fitted  with  a  cork  carrying  two  tubes,  one 
reaching  to  the  bottom,  while  the  other  ends  just  below  the  cork, 
as  shown  in  Fig.  75.  The  former  of  these  tubes  is  connected  to  a 
piece  of  narrow  metal  pipe,  which  is  thrust  into  the  flue  through 

*  Standard  solutions  of  such  a  strength  that  i  c.c.  =  i  c.c.  of  the  gas  to  be 
estimated  (at  N.  P.  T. )  are  sometimes  spoken  of  as  normal  solutions,  or  normal 
gas  solutions.  The  use  of  the  term  normal  solution  for  any  other  solution  than 
those  already  described  on  p.  313,  is  greatly  to  be  deprecated,  as  tending  to 
introduce  confusion. 


Sulplnir  Dioxide  in  Furnace  Gases.  383 

which  the  furnace  gases  are  passing.*  The  other  tube  is  attached 
to  the  large  bottle  A,  filled  with  water,  which  serves  as  an  aspirator. 
The  water  is  allowed  to  flow  slowly  out  of  the  bottle  by  means  of 
the  screw  clamp,  and  is  received  in  a  litre  cylinder.  The  flask 
containing  the  iodine  is  shaken  at  frequent  intervals  during  the 
process. 

When  the  cylinder  is  full  to  the  litre  mark,  it  is  either  changed, 
or  the  clamp  is  closed  for  a  moment  while  the  cylinder  is  emptied, 
and  then  reopened ;  a  record  being  made  of  the  number  of  litres 
thus  drawn  off. 

During  the  experiment  the  temperature  of  the  water  is  taken 
(which  will  be  the  temperature  of  the  gas),  and  the  height  of  the 
barometer  is  noted. 

When  8  or  10  litres  f  of  gas  have  in  this  way  been  aspirated 
through  the  iodine  solution,  the  process  is  stopped,  and  25  c.c.  of 
the  iodine  are  transferred  to  a  beaker  and  titrated  with  the  standard 
thiosulphate.  As  soon  as  the  red-brown  colour  of  the  iodine  solu- 
tion changes  to  a  straw  colour,  one  or  two  drops  of  dilute  starch 
solution  are  added,  and  the  thiosulphate  admitted  drop  by  drop, 
until  the  blue  colour  is  discharged.  A  duplicate  titration  is  made 
in  a  second  portion  of  the  solution. 

EXAMPLE. — Gas  drawn  from  the  flue  of  a  coke  furnace-J 

Volume  of  water  drawn  from  aspirator    ...         ...     8  litres 

Temperature  17° 

Atmospheric  pressure         760  mm. 

Therefore  volume  of  gas  operated  upon  =  — — — —  =  7'53C>7  litres 

at  N.T.P. 

*  This  pipe  should  be  bent  at  a  right  angle  at  a  short  distance  from  the  end, 
and  a  number  of  small  holes  bored  along  the  bent  part,  while  the  extreme  end 
is  closed  either  with  a  metal  plug  or  by  being  roughly  hammered  together.  In 
this  way  the  gas  is  drawn  from  various  points  in  the  horizontal  sectional  area 
of  the  flue. 

f  When  the  gas  is  very  rich  in  sulphur  dioxide,  as,  for  instance,  in  the  flue 
gases  from  the  sulphur-burners  in  vitriol  works,  a  much  smaller  volume,  say  i 
litre,  need  only  be  drawn  through  the  apparatus.  In  such  a  case,  the  pipes 
leading  from  the  absorption  flask  to  the  flue  must  be  filled  with  the  flue  gases 
before  the  tube  is  attached  to  the  flask.  Also,  in  the  final  correction  of  the 
volume  of  gas  operated  upon,  the  volume  of  sulphur  dioxide  which  has  been 
absorbed  must  be  added  to  the  volume  of  water  measured  out  of  the  aspirator. 
In  the  example  here  given,  where  the  amount  of  sulphur  dioxide  is  small, 
these  corrections  would  not  affect  the  third  decimal  figure. 

%  The  furnace  had  been  running  many  hours  previous  to  taking  this  sample, 
and  at  the  time  the  analysis  was  made,  the  draught  of  air  was  checked  by 
partially  closing  the  inlet  at  the  bottom.  These  conditions  were  arranged 
in  order  to  intentionally  send  up  the  percentage  of  sulphur  dioxide  some- 
what. 


3^4  Gas  Analysis. 

100  c.c.  of  iodine  employed  in  the  absorption  flask 

(i  c.c.  =  i  c.c.  thiosulphate  =  o'i  c.c.  SO.2  at  N.T.P.) 

After  absorption,  25  c.c.  iodine  solution  required  17  c.c.  thio- 
sulphate. 

.*.  volume  of  thiosulphate  required  for  100  c.c.  =  17  x  4  =  68  c.c. 

and  volume  of  SO2  absorbed  =  —         — '=  3*2  c.c. 
Hence  75307  c.c.  of  the  furnace  gases  contain  3*2  c.c.  SO2 

or  —         —  =  0*425  percentage  SO2  by  volume 
753°  7 

II.  ESTIMATION  BV  ABSORPTION,  AND  MEASUREMENT  OF  THE 
RESIDUAL  GAS. 

These  processes  differ  from  the  foregoing,  in  that  the"  involve 

(1)  the  manipulation  of  comparatively  small  volumes  of  gas,  and 

(2)  the  accurate  measurement  of  these  volumes. 

For  the  manipulation  of  small  volumes  of  gas  special  appa- 
ratus is  required,  and  for  the  accurate  measurement  of  gaseous 
volumes  special  precautions  are  necessary. 

The  Measurement  of  Gases. — The  gas  whose  volume  is 
to  be  measured  must  be  contained  in  a  graduated  and  accurately 
calibrated  tube,*  and  confined  over  a  suitable  liquid.  Since  no 
gas  is  absolutely  insoluble  in  water,  the  confining  liquid  which  is 
used  in  the  most  exact  processes  of  gas  analysis  is  mercury  ;  but, 
in  the  simpler  forms  of  apparatus  about  to  be  described,  the  con- 
fining liquid  is  water,  which,  with  certain  precautions,  and  in  the 
case  of  a  considerable  number  of  gases,  can  be  employed  with 
fairly  accurate  results. 

The  apparatus  in  which  the  measurements  are  made  is  shown 
in  Fig.  76.  This  is  known  as  a  gas-burette,  and  consists  of  a 
graduated  measuring-tube,  M,  and  a  pressure-tube,  P.  The 
measuring-tube  has  an  ordinary  stop-cock  at  the  top,  and  is  fur- 
nished at  the  bottom  with  a  three-way  tap,  which  allows  of  com- 
munication being  established  with  the  pressure-tube  or  with  the 
outer  air  at  will.  This  will  be  seen  more  clearly  in  Fig.  77.  The 
space  between  the  two  taps  has  a  capacity  of  100  c.c.,  and  is 
graduated  into  cubic  centimetres,  with  subdivisions  of  fifths  of  a 

*  Formerly  these  tubes  were  called  "eudiometers,"  from  the  Greek, 
signifying  to  measure  the  clearness  or  goodness  of  air ;  and  this  term  is  still 
usually  applied  to  such  vessels  when  mercury  is  employed  to  confine  the  gas. 
It  has  become  the  custom,  however,  in  the  more  modern  forms  of  gas  apparatus 
to  call  these  measuring-tubes  ' '  gas-burettes." 


Gas-burettes. 


385 


cubic  centimetre,*  and  the  graduations  are  numbered  in  both 
directions.  Each  of  the  tubes  is  fitted  into  a  loaded  wooden  foot 
or  stand,  cut  in  the  manner  shown  in  the  figure,  so  that  the  tubes 
may  be  brought  as  near  together  as  possible  when  adjusting  the 
level  of  the  liquid  to  the  same  height  in  each.  The  foot  of  the 


FIG.  76. 


FIG.  77. 


measuring-tube  is  made  in  the  form  of  a  clamp,  which  is  shown 
open  in  Fig.  77. 

The  two  tubes  are  connected  by  means  of  a  rubber  tube,  long 
enough  to  allow  one  of  the  tubes  to  be  raised  to  the  top  of  the 

*  The  apparatus  here  figured  is  Hempel's  modified  Winkler's  gas-burette. 
In  the  simple  "  Hempel  "  burette  there  are  no  taps  at  all,  the  top  bein^  closed 
with  a  piece  of  rubber  tube  and  a  pinch-cock. 


386 


Gas  Analysis. 


other.  In  order  to  avoid  taking  this  rubber  tube  off  and  on  the 
burette  when  the  latter  requires  cleaning,  it  is  better  to  introduce 
a  piece  of  glass  tube  as  a  connector,  as  shown  at  /.  When  the 
measuring-tube  and  pressure-tube  have  to  be  disunited,  the  separa- 
tion is  made  by  removing  this  glass  connector. 

The  gas  in  the  measuring-tube  is  confined  over  water,  and  its 
volume  is  measured  at  atmospheric  pressure,  which  is  secured  by 
placing  the  pressure-tube  in  such  a  position  relative  to  the  measuring- 
tube  that  the  water  stands  at  the  same  level  in  each  (Fig.  76). 
When  reading  the  volume  of  the  gas,  the  lower  line  of  meniscus  is 
taken,  just  as  when  reading  the  volume 
of  liquid  in  an  ordinary  burette,  and  the 
aids  may  be  employed  to  render  the 
meniscus  visible,  as  are  described  on  p. 
310.  Before  reading  the  volume  of  a 
gas,  a  regular  interval  of  time  should  be 
allowed  for  the  water  to  drain  down  the 
walls  of  the  burette ;  60  seconds  is  a 
suitable  time. 

Calibration  of  the  Gas-burette, 
Just  as  in  the  case  of  the  ordinary  burette 
used  for  liquids,  the  gas  burette  requires 
to  be  calibrated  before  being  used  for 
analytical  purposes.  This  is  done  in  the 
following  way.  The  measuring-tube  is 
disconnected  from  the  pressure-tube  at 
the  glass  union  /  (Fig.  76),  and  (without 
its  foot)  is  supported  in  an  inverted  posi- 
tion in  an  ordinary  clamp,  as  shown  in 
Fig.  78.  The  tap  a  is  opened,  and  B  is 
turned  so  as  to  communicate  between  the 
burette  and  the  rubber  tube.  A  beaker 
of  water,  at  the  temperature  of  the  room, 
is  brought  under  the  bottom  end  of  the 
pipette,  and  the  liquid  sucked  up  until  the 
burette  is  filled  above  the  tap  B,  when 
the  bottom  tap  is  closed,  and  the  rubber 
removed.  By  cautiously  opening  a  (B 
being  left  open  all  the  time),  water  is 
allowed  to  drop  out  until  the  liquid  is 
just  level  with  the  under  side  of  the  tap  B,  when  a  is  again  closed. 
A  definite  volume  (3,  4,  or  5  c.c.)  of  the  water  is  then  run  out 


FIG.  78. 


Correction  of  Gaseous    Vohimes.  387 

at  a  time,*  and  weighed  exactly  in  the  manner  described  on  p.  309  ; 
and  from  the  weight  so  obtained,  the  true  capacity  of  the  graduations 
upon  the  tube  are  calculated  and  tabulated,  as  in  the  case  of  ordinary 
burettes  for  liquids. 

Correction  of  Gaseous  Volumes. — The  volume  which  a 
given  weight  of  gas  will  occupy  depends  upon  three  conditions, 
namely  (i)  the  temperature,  (2)  the  pressure,  and  (3)  the  degree 
of  humidity  of  the  gas,  at  the  time  the  measurement  is  made. 

In  order,  therefore,  that  the  various  volumes  observed  during  an 
analysis  shall  be  comparable  with  each  other,  it  is  necessary  either 
that  the  conditions  mentioned  should  remain  constant  throughout, 
or  that  the  volumes  measured  under  different,  but  observed,  con- 
ditions, should  be  reduced  by  calculation  to  a  common  standard. 

In  exact  methods  of  analysis,  the  latter  plan  is  invariably 
adopted  ;  but  in  the  more  rapid  and  somewhat  rougher  methods 
employed  for  technical  purposes,  the  analysis  may  usually  be  carried 
out  without  disturbing  the  uniformity  of  the  conditions  to  an  extent 
which  will  introduce  any  material  error  in  the  results. 

The  recognised  standard  to  which  gaseous  volumes  are  reduced, 
is  the  volume  which  the  gas  would  occupy  at  o°  and  under  a  pressure 
of  760  mm.,  when  in  the  dry  state. 

(a)  Temperature  Correction. — The  coefficient  of  expansion  of 
gases  is  taken  as  o'oc>3665,  or  ^^  ;  therefore  the  volume  at  o°  equals 
the  volume  at  /°  divided  by  i  + 


V  V  x  273 

hence  V0  =  -  — ^ —     or    ~ 

i  +  0-003665 1  273  +  t 

where  V0  =  volume  at  o°,  and  V  =  volume  at  /°. 

(1)}  Pressure  Correction. — The  volume  of  a  gas  being  inversely 
as  the  pressure — 

=  VP 
0  ~  760 

where  V  =  the  volume  at  P  pressure  :  or,  making  both  corrections 
together — 

V   -  VP  VP  x  273 

0      760(1  +  0-003665/)  760  x  (273  +  /) 

(c)  Aqueous  Vapour. — The  aqueous  vapour  present  in  a  gas 
exerts  a  pressure  in  opposition  to  the  barometric  pressure  ;  hence 
the  volume  of  a  gas  is  increased  by  the  presence  of  aqueous  vapour. 

*  At  the  narrow  part  of  the  burette,  near  the  tap  a,  each  graduation  should 
be  calibrated. 


388  Gas  Analysis. 

If  the  gas  is  saturated  with  aqueous  vapour,  and  an  excess  of  water, 
however  minute,  be  present,  then  the  pressure  or  tension  of  the 
aqueous  vapour  is  independent  of  change  of  pressure,  varying  only 
with  change  of  temperature.  The  tension  of  aqueous  vapour  has 
been  exactly  determined  for  every  degree  of  temperature  within 
limits,  and  in  the  table  in  the  Appendix  will  be  found  the  tension  or 
pressure,  in  millimetres  of  mercury,  of  the  vapour  of  water  between 
the  temperatures  of  5°  and  25°. 

In  making  the  necessary  correction  for  aqueous  vapour,  there- 
fore, the  number  of  millimetres  of  mercury  representing  the  tension 
of  aqueous  vapour  at  the  particular  temperature  at  which  the  gas 
measurement  is  made,  is  deducted  from  the  barometric  pressure  to 
which  the  gas  is  exposed.  For  example,  suppose  a  volume  of  gas 
is  measured  at  atmospheric  pressure  when  the  barometer  is  at 
760  mm.,  and  the  temperature  is  15°  ;  then,  on  referring  to  the 
table,  the  tension  of  aqueous  vapour  at  15°  is  seen  to  be  127  mm. 
Deducting  this  from  the  barometric  pressure  gives  760  —  12*7 
=  747*3  mm.  as  the  true  pressure. 

If  p  stands  for  the  pressure  due  to  aqueous  vapour,  then  the 
formula  — 

/)  x  273 


V   -  or 

760(1  +  0-003665;)  760  x  (27.3  +  /) 

embraces  all  the  corrections  necessary  to  reduce  a  volume  of  a  gas, 
saturated  with  aqueous  vapour,  to  the  standard  conditions. 

When  gases  are  confined  over  water,  as  in  the  Hempel-Winkler 
burette,  the  condition  of  complete  saturation  with  aqueous  vapour 
is  of  course  always  present.*  With  this  apparatus  also,  the  gas- 
volumes  are  always  read  at  the  atmospheric  pressure  ;  and  as  the 
analytical  operations  are  rapidly  performed,  changes  of  barometric 
pressure  sufficient  to  influence  the  results  need  not  be  anticipated. 

Changes  of  temperature,  however,  must  be  guarded  against  as 
far  as  possible,  and  with  this  object  in  view,  the  apparatus  should 
be  handled  entirely  by  the  wooden  feet,  and  not  by  the  glass 
portions.  In  order  to  ascertain  the  temperature  of  the  gas  and  see 
how  far  it  is  being  maintained  uniform  throughout,  a  simple  and 
convenient  plan  is  to  suspend  a  thermometer  inside  the  pressure- 
tube  P  (Fig.  76)  by  means  of  a  thread,  so  that  it  reaches  nearly 
to  the  bottom,  and  remains  there  during  the  whole  analysis.  As 
the  water  is  continually  being  passed  backwards  and  forwards  from 

*  When  mercury  is  employed  as  the  confining  liquid,  complete  saturation 
of  the  gas  with  aqueous  vapour  is  ensured  by  introducing  a  drop  of  water  into 
the  measuring-tube,  or  eudiometer. 


Correction  of  Gaseous   Volumes.  389 

the  pressure-tube  to  the  measuring-tube,  the  temperature  of  the  gas 
may  be  taken  as  the  same  as  that  of  the  water  over  which  it  is 
confined  ;  and  if  the  temperature  of  the  latter  does  not  materially 
change,  that  of  the  gas  may  be  considered  as  practically  uniform. 

As  stated  above,  when  the  conditions  under  which  gas  measure- 
ments are  made  are  constant,  it  is  not  necessary  to  reduce  the  ob- 
served volumes  to  standard  conditions.  This  will  be  rendered  more 
obvious  by  the  following  concrete  example.* 

The  original  volume  of  gaseous  mixture  in  the  burette  measured 
100  c.c.  at  atmospheric  pressure  (i.e.  when  the  water  was  at  the 
same  level  in  both  tubes,  M,  P  (Fig.  76).  One  constituent,  ,r,  was 
then  removed  by  absorption,  and  the  gas  measured  again.  Its 
volume  now  was  75  c.c.  at  atmospheric  pressure. 

The  temperature  was  15°,  and  the  barometric  pressure  755  mm. 
throughout. 

Then  (i)  without  making  reduction  to  standard  conditions, 
we  get — 

100-^-75  =  25  =  percentage  of  x  in  the  mixture 
(2)  On   reducing  the  two  volumes  by  means  of  the  formula 


V    — we  Eret — 

v°~  760(1  +0-003665/)' 


- t - 

therefore  92*56  —69*43  =  23*13  =  volume  of  x  in  92  56  c.c.  original 
gas 

and   — —  •  A  C~  —  25  =  percentage  of  x  in  the  mixture. 

Since  the  tension  of  aqueous  vapour  is  independent  of  pressure, 
then,  in  the  event  of  any  alteration  of  the  barometric  pressure 
taking  place  during  an  analysis,  it  is  only  necessary  to  make  a 
correction  for  pressure  ;  not  necessarily  by  reducing  #//the  volumes 
to  the  standard,  but  by  reducing  all  to  the  same  pressure  as  any 
one  of  them. 

Thus,  in  the  above  example,  suppose  that  between  the  two 
measurements  the  barometer  fell  from  755  to  750  mm.,  the  tem- 
perature remaining  constant  at  1 5°,  then  the  following  are  the  data: — 

Original  volume  =  100  c.c.  at  15°  and  755  mm. 
After  absorbing  xt  volume  =    75       „      15°     „    750     „ 

*  The  student  who  has  not  had  practice  in  the  reduction  of  gases  to  normal 
conditions,  is  strongly  recommended  to  carefully  work  out  these  calculations 
for  himself. 


39O  Gas  Analysis. 

then  — —   =.  74'5  =  volume  which  the  residual  gas 

"55  would  occupy  if  measured  under  the 

same   conditions    as    the    original 
volume 
hence  100  —  74  5  =  25*5  —  percentage  of.r  in  this  mixture 

If  now,  from  the  above  data,  the  two  volumes  be  reduced  to 
standard  conditions  by  means  of  the  formula,  as  in  the  last 
example,  it  will  be  found  that  the  same  result  is  obtained,  namely, 
25 '5  percentage  of  or. 

Again,  since  the  tension  of  aqueous  vapour  depends  upon  the 
temperature,  increasing  with  rise  of  temperature,  change  of  tempe- 
rature obviously  will  produce  an  alteration  of  the  pressure,  even 
though  the  barometric  pressure  remains  constant.  For  example, 
suppose  in  the  above  illustration  the  100  c.c.  original  volume  is 
measured  at  15°,  and  the  75  c.c.  residual  gas  is  measured  at  20°, 
the  barometer  standing  uniformly  at  760  mm.  ;  then  the  actual 
pressure  in  the  first  case  is  760—  127  (tension  of  aqueous  vapour 
at  15°),  and  in  the  second  it  is  760  —  17*4  (the  tension  at  20°). 

Hence,  if  any  change  of  temperature  is  observed  in  the 
gas  during  the  progress  of  an  analysis,  the  observed  volumes 
must  be  reduced  to  the  standard  by  means  of  the  formula 

=  _       V(P-l)       _  . 
760(1  +  0-0036650  ' 

Collection  of  Gas  for  Analysis.— If  the  gas  for  analysis 
is  collected  in  the  laboratory,  say,  a  sample  of  ordinary  coal-gas, 
it  may  be  introduced  into  the  burette  by  first  placing  the  measuring- 
tube  on  a  higher  level  than  the  pressure-tube,  and  allowing  the 
former  to  empty.  By  means  of  the  three-way  tap,  communication 
between  the  measuring-tube  and  the  outer  air  is  then  opened,  and 
a  rapid  stream  of  the  gas  passed  through  the  tube  from  the  top, 

*  In  the  Appendix  there  will  be  found  a  table  of  factors  for  reducing  volumes 
of  gas  to  standard  temperature  and  pressure,  by  the  use  of  which  much  time 
and  calculation  will  be  saved.  In  the  columns  opposite  to  the  pressure  and 
temperature  at  which  a  gas  has  been  measured,  a  figure  is  found,  which  when 
multiplied  into  the  observed  volume  will  at  once  give  the  corrected  volume. 
Thus,  suppose  60  c.c.  of  gas  saturated  with  aqueous  vapour  were  measured  at 
12°  with  the  barometer  standing  at  766  mm.  Then — 

766  —  10  5  (tension  aq.  vap.  at  12°)  =  755 -5  mm.  =  actual  pressure 

On  referring  to  the  table,  the  figure  in  the  12°  column  which  falls  halfway 
between  755  and  756  mm.  is  0^9522  ; 

therefore  60X0-9522  =  S7'132  =  corrected  volume 

The  student  is  warned  against  using  such  factors  before  he  is  perfectly  familiar 
with  the  method  of  calculating  out  the  correction  for  himself.  Such  a  table  is 
not  intended  to  save  him  the  trouble  of  learning  how  to  make  the  necessary 
calculations,  but  only  to  save  the  time  and  trouble  of  continually  repeating  the 
saoie  rather  lengthy  calculations  when  he  does  know  how  to  make  them. 


Correction  of  Samples  of  Gas. 


391 


until  the  air  has  been  entirely  swept  out.  The  upper  tap  is  then 
closed,  and  the  lower  one  turned  so  as  to  re-establish  communica- 
tion with  the  pressure-tube. 

When  the  available  supply  of  the  gas  is  comparatively  small, 
it  may  be  collected  in  a  glass  tube  over  water  (or,  if  necessary,  over 
mercury),  and  afterwards  transferred  to  the  burette  as  described 
below.  The  tube  may  conveniently  have  the  form  shown  in  Fig. 
79.  It  is  first  filled  with  water  by  sucking  the  liquid  up  and 

t 


FIG.  79. 


FIG.  80. 


closing  the  rubber  with  a  pinch-cock,  and  the  gas  is  then  passed 
up  from  below  in  the  usual  manner.  The  tube  is  then  connected 
to  the  gas  burette,  with  the  same  precautions  against  enclosing  air 
in  the  joints  as  given  below. 

When  the  gas  is  collected  away  from  the  laboratory,  it  should 
be  taken  in  glass  tubes  drawn  out  to  a  capillary  constriction  at 
each  end.  These  tubes  are  filled  either  by  aspirating  the  gns 


392 


Gas  Analysis. 


through  them  so  as  to  sweep  out  the  air,  and  then  hermetically 
sealing  them  at  the  constrictions  ;  or  by  taking  them  to  the  spot 
in  a  vacuous  and  sealed  condition,  and  then  breaking  open  one  end 
in  the  gas  to  be  collected.  After  the  gas  has  filled  the  tube,  the 
end  is  again  hermetically  sealed. 

In  order  to  transfer  the  gas  from  the  sealed  tube  to  the  burette, 
a  piece  of  capillary  tube,  /,  bent  twice  at  right  angles,  is  attached 
to  the  latter,  as  shown  in  Figs.  79  and  80,  the  joints  being  wired 
round.  T.he  pressure-tube  is  then  raised,  until  water  completely 
fills  the  measuring-tube  and  drops  from  the  open  end  of  the  bent 
capillary.* 

Near  to  each  end  of  the  sealed  tube,  a  slight  scratch  with  a  file 
is  made.  Over  one  end  a  short  piece  of  rubber  tube  is  slipped, 
and  the  projecting  portion  of  it  filled  up  with  water.  The  bent 
capillary  (already  entirely  full  of  water)  is  then  introduced  into  this 
tube,  and  the  latter  secured  with  binding  wire.  In  this  way  all  air 
is  excluded  from  the  joint.  The  lower  end  of  the  tube  is  dipped 
into  a  vessel  of  water. 

The  tube   is  broken  at    the  file-mark  within   the  indiarubber 

joint,  and  the  end  beneath  the 
water  is  broken  off  by  means  of  a 
pair  of  pliers.  On  lowering  the 
pressure-tube  and  opening  the  tap 
a  at  the  top  of  the  measuring-tube, 
the  gas  will  be  drawn  over  into 
the  burette. 

Absorption  of  Gases.  — 
Except  when  the  absorbing  liquid 
is  water  (in  which  case  the  absorp- 
tion is  made  in  the  measuring- 
tube  of  the  burette),  the  absorption 
of  the  gases  in  a  mixture  is  carried 
out  in  a  separate  piece  of  apparatus. 
Fig.  8 1  shows  the  Hempel  gas- 
pipette,  which  is  employed  for 
this  purpose.  It  consists  of  two 
bulbs  connected  by  a  bent  tube, 
mounted  upon  a  wooden  stand.  The  capacity  of  the  larger  bulb 
should  be  at  least  150  c.c.,  so  that  when  the  gas  from  the  full 

*  Whenever  the  available  supply  of  the  gas  to  be  analysed  renders  such  a 
course  possible,  the  water  used  in  the  burette  should  be  first  saturated  with  the 
gas  by  shaking  a  quantity  of  water  with  some  of  the  gas  in  a  stoppered  bottle 
for  a  short  time. 


FIG.  81. 


A  bsorption  Pipettes. 


393 


burette  (i.e.  100  c.c.)  is  transferred  to  it,  there  will  be  room  for  a 
considerable  quantity  of  the  reagent  also.  The  bent  tube  d  has  a 
capillary  bore,  and  a  strip  of  white  glass  or  porcelain  is  fixed 
behind  it,  in  order  to  render  the  position  of  the  liquid  in  the  tube 
more  visible.  Upon  the  end  of  this  tube  a  short  piece  of  moderately 
stout  rubber  tube  is  secured  with  wire,  and  the  tube  furnished  with 
a  pinch-cock.  The  pipette  is  filled  by  pouring  the  reagent  through 
the  rather  wide  tube  b  (for  which  purpose  a  thistle-funnel  should 
be  used,  so  as  to  avoid  spilling  the  reagent  over  the  outside),  the 
clamp  c  being  open.  Such  a  volume  of  the  reagent  is  introduced, 
that  when  it  is  sucked  up  into  the  capillary  tube  nearly  to  the  top, 
there  shall  be  just  a  small  quantity  of  the  liquid  in  the  smaller  bulb. 

When  not  in  use,  the  tube  b  should  be  closed  with  a  small  cork, 
the  rubber  tube  being  closed,  not  with  the  pinch-cock  (which  spoils 
the  tube  after  a  time),  but  by  inserting  a  short  piece  of  glass  rod 
into  the  tube. 

A  separate  pipette  is  used  for  each  reagent,  and  they  should  be 
distinctly  labelled  so  as  to  indicate  the  reagent  they  contain. 

In  cases  where  it  is  necessary  to  guard  the  reagent  from  contact 
with  the  atmosphere  (as,  for  example,  with  alkaline  pyrogallol, 
cuprous  chloride,  etc.),  the 
double  pipette  shown  in  Fig.  82 
is  employed.  It  differs  from 
the  simple  pipette  only  by  the 
two  extra  bulbs  b,  b',  which,  by 
being  partially  filled  with  water, 
serve  as  a  guard  or  water-seal. 

This  simple  addition  to  the 
pipette,  however,  makes  it 
somewhat  more  difficult  to  fill. 
It  is  necessary  so  to  arrange 
matters  that  when  the  absorb- 
ing reagent  fills  the  bulb  a, 
the  water  shall  occupy  bulb  £, 
so  that  when  the  reagent  passes 
up  into  «',  the  water  shall  be 
driven  into  b'.  If  this  condi- 
tion is  not  properly  secured,  as  the  reagent  is  made  to  pass  back- 
wards and  forwards  between  a  and  a',  either  air  will  be  drawn  in 
through  the  water  in  the  water-seal,  or  else  some  of  the  water  itself 
will  be  drawn  over  into  a'.  The  following  is  the  best  method  of 
filling  the  apparatus  : — 


FIG.  82. 


394 


Gas  Analysis. 


The  empty  pipette  is  supported  in  an  inverted  position,  and  an 
ordinary  lo-c.c.  pipette  is  connected  to  the  capillary  tube,  in  the 
manner  shown  in  Fig.  83.  To  the  free  end  of  the  latter  a  piece  of 
narrow  glass  tube,  /,  is  attached  by  means  of  a  rubber  union  provided 

with  a  pinch-cock.  A  short 
length  of  rubber  tube,  also 
carrying  a  pinch  -  cock,  is 
attached  to  the  other  end  of 
the  gas-pipette.  The  air  within 
the  apparatus  is  then  swept 
out  by  passing  a  stream  of 
some  inert  gas  (i.e.  inert  to- 
wards the  particular  reagent 
which  is  destined  to  fill  the 
pipette),  such  as  nitrogen,  or 
carbon  dioxide.  The  narrow 
glass  tube  is  then  dipped  into 
the  bottle  containing  the  re- 
agent in  question,  which  is 
drawn  up  into  the  apparatus 
by  applying  suction  through 
the  rubber  tube  s.  As  soon 
as  the  bulb  a  is  completely 
filled  (care  being  taken  not  to 
draw  any  of  the  liquid  over 
into  bttlb  #'),  the  clamps  c  and 
c'  are  closed,  and  the  tube  / 
disconnected.  The  apparatus 
is  then  turned  over  into  its 
normal  position,  the  little 
FIG.  83.  pipette  P  being  supported  by 

the  hand. 

The  burette  (which  has  previously  been  filled  with  the  same 
inert  gas  used  for  the  pipette)  is  now  attached  to  the  rubber  union 
at  c  by  means  of  the  bent  capillary  tube,  as  shown  in  Fig.  84.  The 
reagent  should  now  occupy  the  entire  space  between  the  clamp  c, 
and  a  point  x  in  the  bent  tube,  except  probably  for  a  small  bubble 
which  will  have  collected  at  y. 

The  rubber  tube  S  is  now  removed,  and  3  or  4  c.c.  of  water  are 
introduced  into  the  bulb  b'  by  means  of  a  thistle  funnel.  This 
water  will  partly  descend  into  the  bent  tube  uniting  b1  with  £,  and 
is  intended  to  serve  as  the  temporary  water-seal. 


Filling  the  Double  Pipette. 


395 


The  inert  gas  in  the  burette  is  now  slowly  passed  into  the 
pipette  by  raising  the  pressure-tube  and  opening  the  tap  of  the 
measuring-tube  and  the  clamps  c  and  d.  As  the  gas  passes  in,  the 
reagent  is  driven  from  bulb  a  to  a',  while  the  gas  which  was  in  a 
is  expelled  through  the  small  quantity  of  water  in  b' .  When  the 


FIG.  84. 

whole  of  the  gas  from  the  burette  has  been  transferred  to  the  pipette, 
clamp  c  is  closed,  and  water  is  introduced  into  the  bulb  b'  until  it 
is  nearly  filled. 

The  clamp  d  is  now  closed,  and  the  little  pipette  removed. 
The  clamp  should  be  opened  again  just  to  allow  the  liquids  to  sink 
to  their  natural  level,  and  the  rubber  then  closed  by  means  of  a 


396  Gas  Analysis. 

plug  of  glass  rod.  The  gas-pipette  is  now  properly  charged,  the 
space  between  the  reagent  and  the  water-seal  being  occupied  by 
inert  gas,  while  the  confining  water  occupies  such  a  position  in  the 
bulbs  b  and  b'  that  it  will  neither  pass  over  into  a',  nor  allow  air  to 
pass  through  when  the  reagent  is  being  transferred  backwards  and 
forwards  from  a  to  a'. 

The  interposition  of  the  lo-c.c.  pipette  in  the  filling  operation 
will  have  secured  the  introduction  of  rather  more  than  enough  of 
the  reagent  to  fill  the  bulb  a.  When,  therefore,  in  the  process  of 
returning  a  gas  from  the  absorption-pipette  to  the  burette,  the  re- 
agent completely  fills  the  bulb  a  and  the  capillary  tube,  there  will 
still  remain  a  few  cubic  centimetres  in  bulb  a'. 

Since  100  c.c.  of  gas  (the  capacity  of  the  burette)  may  at  any 
time  be  introduced  into  bulb  a,  it  will  be  evident  that  the  capacity 
of  a'}  by  and  b'  must  not  be  less  than  this  volume,  otherwise  they 
will  overflow.  Badly  blown  specimens  are  sometimes  met  with 
which  have  this  defect  ;  it  will  be  at  once  discovered  during  the 
operation  of  filling,  and  the  apparatus  must  be  rejected. 

Gases  commonly  estimated  by  Direct  Absorption.— 
The  gases  which  are  most  commonly  estimated  by  simple  absorp- 
tion in  a  gas-pipette,  and  the  reagents  that  are  employed  for  the 
purpose  of  their  absorption,  are  the  following  : — 

(1)  Carbon  dioxide;  absorbed  by  potassium  hydroxide. 
The  reagent   is   made   by   dissolving    150  grams   of  commercial 
caustic  potash  in  500  c.c.  of  water,  and  is  used  in  the  simple  two- 
bulb  pipette. 

(2)  Carbon  monoxide;   absorbed    by   cuprous   chloride. 
Cuprous  chloride  is  soluble  both  in  ammonia  and  in  hydrochloric 
acid,  and  a  solution   in  either  solvent  may  be  employed  for  the 
absorption  of  carbon  monoxide  (see  footnote,  p.  86).      The  acid 
solution  is   preferable,  however,  except  under   the  circumstances 
described  on  p.  408,  when  the  ammoniacal  solution  must  be  used. 

(a)  Add  Solution. — Thirty  grams  of  cuprous   chloride*    are 

*  Cuprous  chloride  is  supplied  by  the  dealers  in  a  fairly  pure  state.  Should 
the  student  require  to  prepare  it  for  himself,  it  may  readily  be  obtained  by  dis- 
solving 25  grams  of  copper  oxide  in  aco  c.c.  of  strong  hydrochloric  acid,  and 
adding  about  25  grams  of  copper  clippings  or  wire.  The  mixture  is  then  gently 
boiled  for  about  an  hour,  more  acid  being  added  if  the  solution  becomes  much 
reduced  in  bulk  by  evaporation.  The  dark  solution  is  then  poured  into  a  large 
beaker  of  water,  and  the  white  precipitate  of  cuprous  chloride  allowed  to  settle. 
The  clear  liquid  is  decanted  off,  and  the  precipitate  transferred  to  a  flask  or 
bottle,  where  it  is  again  allowed  to  settle  and  the  liquid  poured  off.  Hydro- 
chloric acid  (strong  acid  3  parts,  water  i  part)  is  then  added  in  sufficient 
quantity  just  to  dissolve  the  cuprous  ohloride.  A  few  strips  of  copper  are 
introduced,  and  the  flask  or  bottle  closed  with  a  good  cork  (a  glass  stopper 
should  not  be  used). 


Gases  Estimated  by  Direct  Absorption.         397 

mixed  with  50  c.c.  of  water  in  a  £-litre  flask,  and  150  c.c.  of  strong 
hydrochloric  acid  added.  If  a  few  copper  turnings  or  thin  strips 
of  the  metal  be  put  into  the  brownish  solution,  and  the  flask  corked 
up  and  left  for  a  day  or  so,  the  liquid  will  become  colourless. 

(b}  Aminoniacal  Solution. — Twenty  grams  of  cuprous  chloride 
are  mixed  with  150  c.c.  of  water  in  a  flask  fitted  with  a  cork  carry- 
ing two  tubes,  one  reaching  to  the  bottom,  while  the  other  ends 
just  below  the  cork.  The  air  is  swept  out  of  the  flask  by  a  stream 
of  indifferent  gas  (hydrogen  or  carbon  dioxide),  after  which  the 
exit  tube  is  made  to  dip  beneath  water.  A  stream  of  ammonia  is 
then  passed  into  the  solution  (obtained  by  gently  heating  a  strong 
solution  of  ammonia  in  a  separate  flask,  the  latter  not  being  con- 
nected until  the  issuing  ammonia  has  expelled  the  air)  until  the 
cuprous  chloride  has  entirely  dissolved.  Any  unnecessary  excess 
of  ammonia  should  be  avoided.  This  reagent  is  used  in  the  double 
pipette,  and  must  be  exposed  to  the  air  as  little  as  possible  during 
the  process  of  filling  the  pipette. 

(3)  Oxygen ;    absorbed    by   alkaline    pyrogallol.*      The 
reagent  is  prepared  by  dissolving  20  grams  of  pyrogallol  in  200  c.c. 
of  the  potassium  hydroxide  solution  described  above  (No.  i). 

This  reagent  is  used  in  the  double  pipette. 

Should  the  sample  of  gas  under  analysis  contain  more  than 
about  25  per  cent,  of  oxygen  (which  is  seldom  the  case),  it  should 
be  diluted  with  a  known  volume  of  hydrogen  until  the  proportion 
of  oxygen  is  reduced  to  within  this  amount  ;  otherwise  the  traces 
of  carbon  monoxide  which  are  liberated  during  the  absorption  of 
oxygen  by  alkaline  pyrogallol  will  prejudice  the  result.  If  the 
original  gas,  however,  is  known  to  contain  no  carbon  monoxide, 
this  difficulty  may  be  obviated  by  transferring  the  gas,  after  absorp- 
tion in  the  pyrogallol,  to  the  pipette  containing  cuprous  chloride. 

(4)  Hydrocarbons    (olifines)  ;    absorbed    by  fuming    sul- 
phuric acid.     This  reagent  is  used  in  the  single  pipette. 

The  same  hydrocarbons  may  also  be  absorbed  by  bromine 
Water,  prepared  by  saturating  water  with  bromine.  After  ex- 
posure to  either  of  these  reagents,  the  gas  is  freed  from  sulphur 
dioxide  or  from  the  vapour  of  bromine  by  being  transferred  to  the 
caustic  potash  pipette.  The  tubes  of  the  pipette  containing  the 
fuming  sulphuric  acid  must  be  closed  by  means  of  a  piece  of  glass 
rod  and  rubber  tube  when  the  apparatus  is  not  in  use. 

Benzene  vapour  is  absorbed  by  fuming  sulphuric  acid,  but  not 
by  bromine  water.     It  is  also  absorbed  by  fuming  nitric  acid. 
*  Sometimes,  but  incorrectly,  called  "  pyrogallic  acid." 


398  Gas  Analysis. 

(5)  Nitric  oxide;   absorbed    by  ferrous   sulphate. — The 

solution  of  ferrous  sulphate   is   prepared   by  dissolving  70  grams 
of  the  salt  in  150  c.c.  of  water.     It  is  used  in  a  double  pipette. 

[Potassium  permanganate  acidified  with  sulphuric  acid  may 
also  be  used  for  the  absorption  of  nitric  oxide ;  in  this  case  the 
simple  pipette  is  employed.] 

(6)  Chlorine,  sulphuretted   hydrogen,    sulphur   dioxide,  hydro- 
chloric  acid,  and   acid  gases   generally,  may  be   absorbed   from 
gaseous  mixtures  by  caustic  potash. 

Estimation  of  Carbon  Dioxide,  Oxygen,  Carbon  Mon- 
oxide and  Nitrogen. — As  a  first  exercise  in  manipulating  the 
gas-apparatus,  an  absorption-analysis  may  be  made  of  an  artificially 
prepared  mixture  of  these  four  gases,  most  of  which  are  very 
constantly  met  with  associated  together  in  such  gaseous  mixtures 
as  furnace  gases,  generator  gases,  water-gas,  coal-gas,  etc. 

The  mixture  may  be  prepared  by  partially  filling  (say  about 
three-fourths)  the  collecting-tube  shown  in  Fig.  79,  p.  391,  with  air. 
A  small  quantity  of  crystallised  oxalic  acid  is  heated  with  strong 
sulphuric  acid  in  a  test-tube  fitted  with  a  cork  and  delivery-tube, 
and  the  mixed  oxides  of  carbon  collected  in  the  tube  so  as  to  fill 
the  remaining  one-fourth.  The  tube  now  contains  the  four  gases 
in  question. 

They  will  be  estimated  by  absorption  in  the  following  order  :  * — 

(1)  Carbon  dioxide  ;  absorbed  by  caustic  potash. 

(2)  Oxygen  ;  absorbed  by  alkaline  pyrogallol. 

(3)  Carbon  monoxide  ;  absorbed  by  cuprous  chloride. 

(4)  Nitrogen  ;  estimated  by  difference. 

(i)  Saturating  the  Water  for  the  Burette.— A  stoppered 
bottle  of  about  300  c.c.  capacity  is  filled  with  water  and  inverted 
in  a  water-trough.  About  100  c.c.  of  the  gaseous  mixture  is 
bubbled  up  into  the  bottle,  which  is  then  closed  with  the  stopper, 
and  the  gas  thoroughly  shaken  up  with  the  remaining  water  for 
a  few  minutes. 

Some  of  this  water  (now  saturated  with  the  gas  to  be  analysed) 
is  poured  into  the  pressure-tube,  and  the  tap  at  the  foot  of  the 
measuring-tube  is  turned  so  as  to  establish  communication  between 
the  two  tubes.  (This  tap  is  not  again  touched  throughout  the 
analysis,  the  passage  of  the  gas  to  and  from  the  measuring-tube 

*  It  will  be  obvious  that  the  order  in  which  the  gases  are  to  be  absorbed 
from  a  mixture  must  be  carefully  considered.  Thus,  in  the  above  illustration 
the  oxygen  must  be  absorbed  before  the  carbon  monoxide,  otherwise  the 
reagent  "used  to  remove  the  latter  gas  would  be  acted  upon  by  the  oxygen 
present  (see  p.  86). 


Typical  Analysis  by  Absorption. 


399 


and  the  various   pipettes  being  controlled  entirely  by  the  upper 
tap.) 

The  bent  capillary  connecting-tube  is  then  attached  to  the  top 
of  the  measuring-tube,  and  the  latter  is  completely  filled  with  water 
by  raising  the  pressure-tube  and  opening  the  tap  at  the  top,  until 
the  liquid  drips  from  the  end  of  the  capillary  tube.  The  tap  is 
then  closed.  The  indiarubber  connector  upon  the  collecting-tube 
containing  the  gas  is  filled  up  with  a  drop  of  water,  and  then  joined 
to  the  end  of  the  bent  capillary  (see  Fig.  79,  p.  391),  and  the  con- 
nections secured  with  wire.  The  pressure-tube  is  lowered,  and 
the  pinch-cock  upon  the  collecting-tube  and  the  upper  tap  of  the 
measuring-tube  are  opened,  whereby  gas  is  drawn  over  into  the 
latter  tube.  When  sufficient  gas  has  thus  been  transferred,  the  pinch- 
cock  and  tap  are  closed,  and  the  two  tubes  disconnected.  One 
minute  is  allowed  to  elapse  for  the  water  to  drain  down  the  walls 
of  the  measuring-tube,  when  the  volume  of  the  gas  introduced  is 
read  off  by  lowering  the  pres- 
sure-tube until  the  level  of  the 
water  in  it  and  in  the  measur- 
ing-tube is  the  same.  The 
graduation  mark  which  coin^ 
cides  with  the  bottom  of  the 
meniscus  represents  the  volume 
of  gas  taken  for  the  analysis.* 

(2)  Estimation  of  Car- 
bon Dioxide. — The  burette 
is  now  attached  to  the  absorp- 
tion pipette  containing  caustic 
potash  in  the  manner  shown  in 
Fig.  85.  Before  the  two  pieces 
of  apparatus  are  thus  united, 
the  potash  solution  is  drawn 
up  so  as  to  completely  fill  the 
bulb  a,  and  the  bent  capillary 
tube  to  the  mark  c,  which  is  FlG-  85- 

made  upon   the  white   tablet.     The   pinch-cock  keeps  it  in  this 
position.     After  the  rubber  connections  have  been  secured  with 

*  It  is  convenient,  when  possible,  to  employ  100  c.c.  of  the  gas  under 
analysis,  in  which  case  the  number  of  cubic  centimetres  of  the  various  consti- 
tuents which  are  absorbed,  represents  the  percentage  of  each  ingredient  in  the 
gas  mixture.  If,  therefore,  more  than  100  c.c.  has  been  first  introduced  into 
the  apparatus,  the  excess  may  be  removed  by  raising  the  pressure-tube  until 
the  gas  is  compressed  to  exactly  100  c.c.  ;  then,  keeping  the  water  in  that 


4OO  Gas  Analysis. 

wire,  the  pinch-cock  is  opened,  when,  if  the  joints  are  tight,  the 
reagent  will  not  sink  from  its  position  at  c.  The  pinch-cock  may 
be  kept  open  by  raising  it  so  that  it  nips  the  glass  tube. 

The  pressure-tube  is  now  raised  (being  handled  exclusively  by 
the  foot),  and  the  tap  ^/at  tha  top  of  the  measuring-tube  is  opened. 
The  gas  is  thus  transferred  completely  to  the  bulb  a  of  the  pipette. 
Two  or  three  drops  of  water  from  the  measuring-tube  are  allowed 
to  follow  the  gas  into  the  bulb  (by  so  doing  the  capillary  tube  of 
the  pipette  is  washed  each  time  the  apparatus  is  used),  after  which 
the  tap  d  is  closed. 

The  gas  is  allowed  to  remain  in  contact  with  the  potash  for 
about  five  minutes,  during  which  time  the  apparatus  is  gently 
shaken  so  as  to  moisten  the  sides  of  the  bulb  with  the  reagent. 

The  pressure-tube  is  then  lowered,  and  tap  d  opened,  whereby 
the  gas  is  returned  to  the  measuring-tube.  As  soon  as  the  potash 
reaches .  the  point  in  the  capillary  tube  opposite  the  mark  c,  the 
tap  is  closed. 

On  no  account  must  the  reagent  be  allowed  to  reach  the  rubber 
connection,  or  to  pass  over  into  the  measuring-tube. 

The  pressure-tube  is  then  held  in  such  a  position  that  the  water 
stands  at  the  same  level  as  in  the  measuring-tube,  and,  after  waiting 
a  minute  for  the  water  to  drain  down  the  walls  of  the  tube,  the 
volume  is  read  off. 

In  order  to  be  sure  that  the  absorption  of  the  carbon  dioxide 
has  been  complete,  the  gas  is  transferred  once  more  to  the  pipette, 
and  gently  shaken  with  the  potash  as  before,  after  which  it  is 
again  returned  to  the  measuring-tube,  and  its  volume  read  off  at 
atmospheric  pressure,  allowing  the  same  interval  for  the  water  to 
run  down.  If  the  two  readings  agree,  the  first  absorption  was 
complete.* 

Original  volume  of  gas  100  c.c. 

Volume  after  absorption  by  potash  ...    88    ,, 

Carbon  dioxide        12    „   =  12  per  cent. 

position  by  pressing  a  finger  upon  the  rubber  tube,  the  tap  at  the  top  of  the 
measuring-tube  is  momentarily  opened.  This  lets  the  excess  of  gas  escape, 
leaving  100  c.c.  at  atmospheric  pressure.  This  is  controlled  by  again  lowering 
the  pressure-tube  until  the  water  in  each  tube  is  at  the  same  level,  when  the 
gas  should  be  found  to  occupy  100  c.c. 

*  If  due  care  be  taken  to  avoid  handling  the  apparatus,  except  by  the  foot 
of  the  pressure-tube,  the  temperature  of  the  gas  will  not  undergo  any  material 
alteration  during  the  course  of  an  analysis.  The  temperature  may  be  watched 
by  suspending  a  thermometer  in  the  water  in  the  pressure-tube,  as  explained 
on  p.  388. 


Typical  Analysis  by  Absorption.  401 

(3)  Estimation  of  the  Oxygen.— The   potash   pipette  is 
detached   from  the  bent  capillary  tube  at  the  joint  immediately 
above  the  pinch-cock,  Fig.  85,  and  replaced  by  the  double  pipette 
containing  the  alkaline   solution  of  pyrogallol.      Before  the  latter 
is  connected,  the  reagent  is  drawn  up  into  the  capillary  tube  to  a 
marked  point,  which  is  as  nearly  as  possible  the   same  distance 
from  the  pinch-cock  as  that  upon  the  previous  pipette.     The  gas  is 
transferred  to  the  pipette  for  absorption  exactly  as  in  the  former 
case,  and  is  left  in  contact  with  the  reagent,  with  occasional  gentle 
shaking,  for  ten  minutes.     It  is  then  returned  to  the  measuring- 
tube,  the  same  care  being  taken  to  bring  the  reagent  exactly  to 
the  mark  upon  the  capillary  tube.     The  tap  is  closed,  and,  after 
allowing  time  for  the  water  to  drain  off  the  walls  of  the  measuring- 
tube,  the  volume  of  the  residual  gas  is  read  off  at  atmospheric 
pressure. 

Volume  of  gas  before  absorption  of  oxygen     88  c.c. 
,,         »       after  .,  „  73   » 

Oxygen          15   „       ISpercent. 

(4)  Estimation  of  the  Carbon  Monoxide. — The  pyrogallol 
pipette  is  disconnected  and  replaced  by  the  one  containing  cuprous 
chloride  in  acid  solution,  the  reagent  being  previously  drawn  over 
into  the  capillary  tube  to  a  mark  in  the  same  relative  position  as  in 
the  previous  cases. 

The  gas  is  transferred  to  the  pipette  and  allowed  to  remain  ex- 
posed to  the  reagent  for  ten  minutes  ;  after  which  it  is  returned  to 
the  measuring-tube,  and  its   volume   determined  with  the   same 
precautions  as  before. 
Volume  before  absorption  of  carbon 

monoxide 73*0  c.c. 

Volume  of  residual  gas  (nitrogen)     ...     6o'5    „ 

Carbon  monoxide  absorbed    ...     12*5    „    =  12 "5  per  cent. 

Hence  the  composition  of  the  mixture  under  analysis  is- 

Carbon  dioxide      12*0 

Carbon  monoxide 12*5 

Oxygen       15*0 

Nitrogen  (by  difference) 60*5 

100*0 

A  second  analysis  of  the  same  mixture  should  be  made. 


402  Gas  Analysts. 

III.  ESTIMATION  BY  COMBUSTION. 

Hydrogen. — When  hydrogen  and  oxygen  combine,  according 
to  the  equation — 

2H2  +  O2  =  2H2O 

the  two  volumes  of  hydrogen  and  one  volume  of  oxygen  practically 
cease  to  occupy  space,  since  the  volume  of  the  condensed  water  is 
inappreciable.  By  measuring  the  shrinkage  or  contraction  which 
takes  place  under  these  circumstances,  and  multiplying  this  by  §, 
the  volume  of  hydrogen  which  was  burnt  is  ascertained. 

In  order,  therefore,  to  estimate  hydrogen,  a  measured  volume  of 
the  gas  is  mixed  with  a  measured  volume  of  air — rather  greater 
than  is  required  to  furnish  the  necessary  amount  of  oxygen — and 
the  mixture  of  hydrogen  and  oxygen  caused  to  unite  by  one  of  the 
methods  described  below.  After  the  combination  the  residual  gas 
is  measured,  when  two-thirds  of  the  contraction  will  represent  tire 
volume  of  the  hydrogen  consumed. 

Two  methods  are  here  given  for  bringing  about  the  combustion 
of  the  mixture  of  oxygen  and  hydrogen. 

(a)  Combustion  of  Hydrogen  by  means  of  Palladium 
ised  Asbestos. 

Epitome  of  Process.— The  hydrogen  is  mixed  with  the 
necessary  excess  of  air  in  a  gas-burette,  which  is  attached  to  an 
absorption-pipette  charged  only  with  water.  The  capillary  attach- 
ment between  the  two  pieces  of  apparatus  contains  a  thread  of 
asbestos,  upon  which  has  been  deposited  a  quantity  of  finely  divided 
palladium.  As  the  gas  is  transferred  slowly  from  the  burette  to  the 
pipette,  it  passes  over  this  palladiumised  asbestos  (which  is  gently 
warmed  by  a  small  flame),  and  the  oxygen  and  hydrogen  are  thereby 
caused  to  unite.  The  gas  is  finally  returned  to  the  burette  and 
measured. 

As  an  exercise  in  the  process,  a  mixture  of  pure  hydrogen  and 
air  may  be  employed.  About  20  c.c.  of  hydrogen,  as  pure  as 
possible,  are  introduced  into  the  gas-burette,  and  the  volume 
measured.  The  pressure-tube  is  then  lowered,  and  a  quantity  of 
air  drawn  into  the  apparatus  until  the  total  volume  is  about 
80  or  85  c.c. ;  that  is  to  say,  from  60  to  65  c.c.  of  air  are  introduced, 
and  the  volume  is  again  exactly  measured. 

The  gas-burette  is  then  attached  to  a  gas- pipette  charged  with 
water,  by  means  of  the  capillary  combustion-tube  containing  the 
palladiumised  asbestos  *  instead  of  the  usual  connecting-tube.  The 

*  The  little  tube  is  prepared  in  the  following  manner  :  0-25  gram,  of  palla- 
dium foil  is  dissolved  in  the  minimum  quantity  of  aqua  regia  in  a  small  porcelain 
dish,  and  the  solution  evaporated  to  dryness  on  a  -steam-bath.  The  residue  is 


Hydrogen  by  Combustion.  403 

palladium  is  gently  heated  by  brushing  a  Bunsen  flame  along  the 
tube,  which  must  be  kept  warm  during  the  whole  operation,  so  as 
to  prevent  water  (the  product  of  the  combustion)  from  condensing 
in  it.  The  temperature  of  the  tube  must  not  approach  a  visible 
redness. 

The  gas  is  allowed  to  pass  slowly  over  the  warm  palladium, 
which  will  be  seen  to  glow  at  the  end  towards  the  incoming  gas. 
When  the  whole  of  the  gas  has  been  transferred  to  the  pipette,  it  is 
drawn  back  again  into  the  burette.  This  process  is  repeated  once 
or  twice  (if  the  palladiumised  asbestos  is  in  good  order,  one 
repetition  of  the  operation  usually  completes  the  combustion),  after 
which  the  residual  gas  is  measured.  It  is  then  passed  once  more 
into  the  pipette  and  back,  and  measured  again.  If  the  two 
measurements  agree,  the  process  is  complete. 

EXAMPLE. — Original    volume    of   hydrogen    (by    electrolysis) 

—  20'5  C.C. 

Excess  of  air  added.   Total  volume  of  mixture  ... 
Volume  after  combustion 


Therefore  contraction        30*6   „ 

30*6  x  |  =  20*4  c.c.  =  volume  of  hydrogen  found 

moistened  with  three  or  four  drops  of  strong  hydrochloric  acid,  added  from  a 
dropping-tube,  and  20  drops  (i  c.c.)  of  water  added.  The  mixture  maybe 
gently  warmed  to  complete  the  solution.  To  this  red-brown  solution,  when 
cold.  20  drops  of  a  cold  saturated  solution  of  sodium  formate  are  added. 

Into  this  mixture,  which  will  not  exceed  2 '5  c.c.  in  volume,  ©'25  gram  of 
asbestos  thread  is  immersed,  which  will  soak  up  the  whole  of  the  liquid. 

[This  thread  must  be  sufficiently  fine  to  admit  of  being  pushed  into  a 
capillary  tube  of  i  mm.  bore.  It  may  be  obtained  by  unravelling  a  piece  of 
asbestos  cloth  so  as  to  get  single  strands.  It  is  cleansed  from  grease  by 
treating  it  once  or  twice  with  a  little  warm  carbon  disulphide  in  a  test-tube,  and 
then  spreading  it  out  on  a  clean  piece  of  paper  to  dry  ;  after  which  it  should  be 
heated  for  a  few  minutes  on  a  piece  of  platinum  foil.  A  quarter  of  a  gram  of 
this  thread  will  be  about  60  cms.  in  length.] 

A  strong  solution  of  sodium  carbonate  is  added  by  means  of  a  dropping-tube, 
and  gently  worked  into  the  soaked  thread  with  a  glass  rod,  until  the  mixture  is 
alkaline,  and  the  dish  placed  upon  a  steam-bath.  A  gentle  heat  suffices  to 
reduce  the  palladium,  which  is  then  precipitated  throughout  the  asbestos  as  a 
black  deposit.  When  the  contents  of  the  dish  are  dry,  they  are  rinsed  three  or 
four  times  with  hot  water  in  order  to  dissolve  out  the  soluble  salts.  The  thread 
is  then  removed,  and  cut  into  short  pieces  about  4  cms.  long.  One  of  these 
pieces  is  straightened  out  with  a  gentle  twisting  of  the  fingers,  and  laid  upon  a 
piece  of  blotting-paper  for  a  few  minutes  to  remove  the  superfluous  water.  It 
is  then  introduced  into  a  piece  of  thick-walled  capillary  tube,  i  mm.  bore 
and  about  15  cms.  long.  When  the  thread  has  been  pushed  a  little  way  into 
the  tube,  it  may  readily  be  drawn  into  the  middle  by  applying  a  gentle  suction 
to  the  other  end.  The  thread  is  then  dried  by  gently  warming  the  tube  and 
slowly  drawing  air  through  it,  after  which  the  tube  is  bent  at  right  angles, 
about  3  cms.  from  each  end.  The  same  piece  of  palladiumised  asbestos  may 
be  used  for  several  combustions. 


404 


Gas  Analysis. 


(b]  Combustion  of  Hydrogen  by  Explosion  with  Air.— 

For  this  purpose  the  mixture  of  hydrogen  with  excess  of  air  is 
transferred  to  a  special  "  explosion-pipette,"  in  which  are  sealed  two 
platinum  wires,  whereby  the  gaseous  mixture  may  be  ignited  by  a 
spark  from  a  Ruhmkorf  coil. 

Fig.  86  shows  a  form  of  explosion-pipette  in  which  the  gas  is 
confined  over  water.*  Water  which  has  been  acidulated  with 

sulphuric  acid,  and  boiled  to 
expel  dissolved  gases,  is  intro- 
duced at  a  until  the  bulb  b  is 
just  full,  and  the  liquid  stands 
level  in  the  other  limb.  Upon 
the  tubes  a  and  *,  pieces  of  thick- 
walled  rubber  tube  are  securely 
wired,  and  a  pinch-cock  is  placed 
on  each. 

At  d  two  platinum  wires  are 
sealed  into  the  glass,  between 
which  the  electric  spark  is  passed 
when  the  gas  is  to  be  fired.f 
In  the  lower  part  of  the  tube,  at 
c,  two  platinum  electrodes  are 
fused  into  the  glass.  These  are 
for  the  purpose  of  adding  a  small 
quantity  of  "  electrolytic  gas " 
to  the  mixture,  when  the  pro- 
portion of  combustible  gas  is 
so  small  that  no  explosion  will 
take  place  when  the  electric 
spark  is  passed.  Before  the 
"  electrolytic  gas  "  is  generated, 
the  mixture  under  analysis  is 
transferred  to  the  measuring- 
tube,  which  is  then  detached 
from  the  "  explosion-pipette."  The  two  wires  from  the  battery  (not 
from  the  coil)  are  then  connected  to  the  electrodes  (c\  and  the 
oxygen  and  hydrogen  which  are  evolved  are  allowed  to  escape. 

*  Another  form  of  explosion-pipette,  in  which  mercury  is  the  confining 
liquid,  is  shown  and  described  on  p.  407. 

f  To  prevent  the  loops  of  these  thin  wires  from  being  broken  off,  it  is  well  to 
attach  to  each  a  short  piece  of  stouter  platinum  wire,  and  pass  these  through 
holes  in  the  wooden  support  so  that  their  ends  project  and  can  be  bent  ovef 
into  loops,  as  shown  in  the  figure.  The  same  may  be  done  for  the  wires  at  c. 


FIG.  86. 


Hydrogen  by  Explosion.  405 

The  current  is  allowed  to  pass  for  about  fifteen  minutes,  in  order 
to  saturate  the  water,  after  which  the  current  is  stopped  and  the 
liquid  driven  up  to  the  usual  mark  upon  the  capillary.  The  burette 
is  reconnected,  and  the  gas  returned  to  the  pipette.  A  small  quan- 
tity of  electrolytic  gas  is  then  generated  (the  amount  depending 
upon  circumstances),  and  thoroughly  mixed  with  the  gas  already 
present,  before  exploding.  It  is  not  necessary  to  know  the  volume 
of  the  gas  thus  added,  since  it  entirely  disappears  when  fired. 

As  a  first  exercise,  a  mixture  of  pure  hydrogen  and  air  may  be 
exploded.  About  10  to  15  c.c.  of  hydrogen  are  introduced  into  the 
gas-burette,  and  after  being  exactly  measured,  about  60  to  7°  c-c- 
of  air  are  added,  and  the  mixed  gases  again  measured.  The  burette 
is  then  attached  to  the  explosion-pipette,  the  liquid  in  the  latter 
being  previously  drawn  up  to  a  mark  upon  the  capillary  tube  e. 

The  gas  is  then  passed  over  into  the  pipette,  and  the  clamps 
upon  the  rubber  tubes  both  closed.  The  wires  from  the  induction 
coil  are  attached  to  the  wires  at  d,  and  the  electric  spark  allowed 
to  pass.  The  explosion,  although  not  at  all  violent,  will  cause  a 
momentary  expansion  within  the  apparatus,  but  if  sufficient  liquid 
is  present,  no  gas  will  be  driven  out  of  the  bulb-tube.  The  thick 
rubber  tube  upon  a  being  closed,  the  small  quantity  of  air  which 
is  enclosed  in  the  bulb  b  serves  as  a  cushion  or  "  buffer "  at  the 
moment  of  explosion,  and  thus  relieves  the  other  part  of  the 
apparatus  from  undue  pressure. 

The  moment  after  passing  the  spark,  the  tube  upon  a  is  opened  ; 
and  then  the  gas  is  returned  to  the  burette  and  measured.  The 
contraction  represents  the  hydrogen  and  the  atmospheric  oxygen 
with  which  it  has  combined  to  form  water,  and  two-thirds  of  this 
shrinkage  is  the  volume  of  hydrogen  which  was  present. 

EXAMPLE. — Volume  of  hydrogen  taken  =  12*5  c.c. 

Volume  of  hydrogen  and  air       75*8  c.c. 

Volume  after  explosion 57'5    „ 

Contraction  18*3    „ 

i8'3  x  §  =  12*2  c.c.  =  volume  of  hydrogen  found 

For  practice  in  the  use  of  "  electrolytic  gas,"  the  following  ex- 
periment may  be  made  :  A  quantity  of  air  (about  60  c.c.)  is  intro- 
duced into  the  burette,  and  its  volume  measured  in  the  usual  way. 

The  current  from  four  or  five  Grove's  cells  (or  its  equivalent  from 
any  other  convenient  source)  is  passed  through  the  dilute  acid  in 
the  explosion-pipette  by  means  of  the  electrodes  at  c  (Fig.  86)  for 


406  Gas  Analysis. 

about  ten  minutes,  and  the  liquid  shaken  up  once  or  twice  with 
the  gas,  in  order  that  it  may  become  saturated.  The  current  is 
then  interrupted,  and  the  liquid  drawn  up  to  the  majk  upon  the 
capillary.  The  burette  containing  the  measured  volume  of  air  is 
then  attached  in  the  usual  manner.  About  12  c.c.  cf  electrolytic 
gas  are  then  generated  in  the  pipette,  and  drawn  over  into  the 
burette.*  The  mixture  of  air  and  electrolytic  gas  is  then  passed 
twice  backward  and  forward  from  the  burette  to  the  pipette,  in 
order  to  ensure  their  complete  admixture,  after  which  the  clamps 
are  closed,  and  the  mixture  fired  The  residual  gas  is  transferred 
to  the  burette  and  measured,  when  the  volume  should  be  the  same 
as  that  of  the  air  originally  introduced.! 

Methane  (Marsh-gas). — When  a  mixture  of  methane  and  air, 
or  oxygen,  is  exploded,  carbon  dioxide  and  water  are  formed  ;  and 
from  the  following  equation  : — 

CH4  +  2O2  =  CO2  +  2H,O 

it  will  be  seen  that  methane,  when  burnt,  gives  its  own  volume  of 
carbon  dioxide  ;  and  also  that  three  volumes  of  mixed  gases  shrink 
to  one  volume  (i.e.  the  i  volume  of  CO2),  since  the  water  ceases  to 
occupy  space.  The  contraction,  therefore,  is  two-thirds  of  the  total 
volume  of  the  reacting  gases  ;  or,  in  other  words,  the  contraction 
is  equal  to  twice  the  carbon  dioxide  produced,  or  twice  the  volume 
of  the  marsh-gas  burnt. 

Owing  to  the  effect  of  pressure  in  increasing  the  solubility  of 
carbon  dioxide  in  water,  it  is  only  possible  to  obtain  accurate 
results  when  mercury  is  used  as  the  confining  liquid.t  An  explosion- 

*  The  actual  amount  of  electrolytic  gas  which  has  been  added  may  be 
ascertained  by  measuring  the  total  volume  of  gas  now  in  the  burette.  It  is  not 
necessary,  however,  to  know  this  volume  exactly,  and  a  very  little  experience 
will  enable  the  operator  to  judge  of  the  volume  by  the  space  it  occupies  in  the 
explosion-pipette  as  it  is  generated. 

f  Usually  after  the  first  experiment  the  volume  of  the  residual  air  is  not 
quite  identical  with  that  which  was  originally  taken,  owing  lo  the  imperfect 
saturation  of  the  liquid  with  the  various  gases.  If  this  is  the  case,  a  similar 
quantity  of  electrolytic  gas  should  be  again  added,  and  the  mixture  fired  once 
more  after  thorough  admixture.  The  volume  of  the  residual  gas  after  this 
second  explosion  should  then  exactly  agree  with  that  which  was  measured  after 
the  first  operation.  The  process  may  be  repeated  with  varying  amounts  of 
electrolytic  gas,  and  the  volume  of  the  residue  will  be  found  to  remain  constaat. 

£  With  a  view  to  ascertain  the  extent  of  the  loss  of  carbon  dioxide  by 
solution  in  water  when  the  explosion  is  conducted  over  that  liquid,  the  following 
experiments  were  made  with  the  apparatus  figured  on  p.  404  : — 

(i)  67  c.c.  of  a  mixture  of  air  and  carbon  dioxide,  containing  9  per  cent,  of 
CO2>  were  introduced  into  the  apparatus,  and  12  c.c.  of  electrolytic 
gas  added.  After  explosion,  the  volume  remaining  occupied  667  c.c. 
Loss  of  CO2  =  o'3  c.c.  =  0-4  per  cent. 


Marsh- Gas  by  Explosion. 


407 


pipette  in  which  mercury  is  employed  is  shown  in  Fig.  87.  It 
differs  from  the  ordinary  absorption-pipette  only  in  containing  two 
platinum  wires,  fused  into  the  upper 
part  of  bulb  a,  and  in  being  fur- 
nished with  a  stopcock,  in  order 
to  close  the  communication  between 
the  two  bulbs. 

Before  the  apparatus  is  con- 
nected to  the  burette  containing 
the  measured  mixture  for  explosion, 
the  mercury  is  driven  over  into 
the  capillary  tube,  to  a  fixed  mark,* 
by  blowing  through  the  rubber  tube 
upon  the  upper  bulb ;  and  similarly, 
when  the  gas  is  transferred  from 
the  burette  to  the  explosion^pipette, 
it  must  be  drawn  over  by  applying 
suction  to  the  same  rubber  tube. 
The  water  from  the  burette  should 

be  made  to  follow  the  gas  so  as  to  just  fill  the  capillary  tube, 
but  without  allowing  any  to  enter  the  bulb.  Before  firing  the 
mixture,  the  pinch-cock  and  the  tap  are  closed. 


(2)  78  c.c.  of  a  similar  mixture,  containing  20  per  cent.  CO2 ;  15  c.c.  electro- 

lytic gas  added.     Volume  after  explosion  —  77-2  c.c. 

Loss  of  CO2  =  o'8  c.c.  =  i'o  per  cent. 

(3)  5° '8  c.c.  of  similar  mixture,  containing  40  per  cent.  CO2  ;  15  c.c.  elec- 

trolytic gas  added.     Volume  after  explosion  =  49  c.c. 

Loss  of  CO2  =  I'S  c.c.  =  3-3  per  cent. 

About  the  same  volume  of  electrolytic  gas  was  added  three  times,  and  the 
mixture  exploded  and  measured  after  each  addition  ;  the  volumes  obtained 
were  47*4  c.c.,  46-2  c.c.,  and  45^2  c.c.,  showing  a  fairly  regular  loss  of  carbon 
dioxide. 

(4)  47'2  c.c.  of  air,  containing  36-4  per  cent.  CO2 

(a)  20  c.c.  electrolytic  gas  added  ;  after  explosion,  volume  =  45^6  c.c. 
(&}  12        ,,  „  ,,  ex ploded  again,  volume  =  45 x>    ,, 

(c)  20       ,,  ,,  ,,  ,,  „      volume  =  43'6    ,, 

(a)  loss  =1*6  c.c. ,  or  3*4  per  cent. 

(/>)  loss  =  0-6    ,, 

(c)   loss  =  1-4    ,, 

Experiments  i,  2,  and  3  show  that,  with  about  the  same  force  of  explosion, 
the  loss  of  CO2  increases  as  the  percentage  present  rises  ;  while  Experiment  4 
shows  that,  with  the  same  percentage  of  carbon  dioxide,  the  amount  absorbed 
depends  upon  the  force  of  the  explosion. 

*  In  the  apparatus  here  figured,  the  capillary  tube  is  enamelled  white  on 
the  back,  like  the  stem  of  a  thermometer ;  in  other  specimens  the  tube  is  of 
clear  glass,  with  a  white  tablet  fixed  behind,  as  in  Figs.  81,  82. 


408  Gas  Analysis. 

After  the  explosion,  the  gas  is  transferred  to  the  burette  and 
measured.  The  explosion-pipette  is  then  disconnected  and  replaced 
by  the  simple  absorption-pipette  containing  caustic  potash,  and  the 
carbon  dioxide  absorbed  in  the  usual  way. 

Mixtures  of  Hydrogen,  Methane,  and  Nitrogen. 
Many  gaseous  mixtures  which  constantly  come  under  analysis 
(such  as  coal-gas,  producer  gas,  water-gas,  blast-furnace  gases) 
contain  varying  quantities  of  these  three  gases  along  with  others. 
After  all  the  other  gases  have  been  estimated  by  absorption  in  their 
respective  reagents,  the  hydrogen  and  marsh-gas  in  the  residue 
are  determined  by  one  of  the  two  following  methods,  while  the 
nitrogen  is  estimated  by  difference  : — 

(a)  The  gas  is  mixed  with  an  excess  of  air,  and  the  hydrogen 
estimated  by  combustion  by  means  of  palladiumised  asbestos  (p. 
402).  Under  these  conditions,  the  marsh-gas  does  not  burn. 

The  marsh-gas  is  then  determined  by  exploding  the  residual 
mixture  and  absorbing  the  carbon  dioxide,  as  described  above. 

The  volume  of  nitrogen  is  found  by  deducting  from  the  original 
volume  of  gas  the  hydrogen  and  marsh-gas  thus  determined. 

EXAMPLE. — One  hundred  cubic  centimetres  of  coal  gas  were 
exposed  to  the  action  of  the  following  reagents — 

(a)  Caustic  potash  ;  to  absorb  carbon  dioxide. 

(b)  Alkaline  pyrogallol  ;  to  absorb  oxygen. 

(c)  Fuming  sulphuric  acid  ;  to  absorb  olefines  and  benzene 

vapour. 

(d)  Ammoniacal  cuprous  chloride  ;  to  absorb  carbon  mon- 

oxide.* 

The  residual  gas,  measuring  86-6  c.c.,  was  returned  to  the 
cuprous  chloride  pipette,  while  the  burette  was  disconnected  and 
the  water  in  it  (previously  saturated  with  coal-gas)  was  replaced  by 
water  saturated  with  air. 

Twenty  cubic  centimetres  of  the  gas  were  transferred  to  the 
burette  (the  rest  being  reserved  for  a  subsequent  experiment),  and 
air  added  in  more  than  sufficient  quantity  for  the  complete  com- 
bustion of  the  hydrogen. 

Volume  of  gas  =  2o-o  c.c. 
volume  of  gas  +  air  =  64/4   „ 

This   mixture   was  then    passed  over   the  palladiumised  asbestos 
(p.  402). 

*  When  the  absorption  of  carbon  monoxide  is  to  be  followed  by  the  com- 
bustion of  hydrogen  with  palladiumised  asbestos,  the  ammoniacal  solution  of 
cuprous  chloride  should  be  used. 


Mixtures  of  Hydrogen,  Methane,  and  Nitrogen.  409 

Volume  after  combustion  of  hydrogen  =  48*2  c.c. 

therefore  64-4  —  48-2  =  16-2    „    =  contraction 

and  1 6*2  x  §  =  1O8  =  volume  of  hydro- 
gen in  2O'o  c.c.  of  gas 

To  the  residual  gas  (consisting  of  marsh-gas,  nitrogen,  and  a  small 
surplus  of  oxygen)  an  excess  of  oxygen  was  added. 

Volume  of  residual  gas  =  48*2  c.c. 
volume  of  residual  gas  +  oxygen  =  69^6   „ 

The  mixture  was  then  exploded,  and  the  carbon  dioxide  absorbed. 
Volume  after  explosion  =  52*8  c.c. 

therefore  69*6  —  52*8  =  16*8  =  contraction 
volume  after  absorption  of  CO2  =  44-4 

therefore  52*8  —  44*4  =  8*4  c.c.  =  volume   of  CO2   pro- 
duced 
and  therefore  8'4  c.c.  =  volume  of  marsh-gas  in  20  c.c. 

of  gas 

20*0  c.c.  —  (io'8  +  8-4)  =  o'8  c.c.  =  volume  of  nitrogen  in 
20  c.c. 

Since  the  original  volume  of  gas  taken  for  analysis  was  100  c.c. — 

10-8  x  86'6 
then =  46-46  =  per  cent,  of  hydrogen 

and  —  —  as  36-07  =  per  cent,  of  marsh-gas 

,  0-8  x  86'6  c    .. 

and =  3*46  =  per  cent,  of  nitrogen 

(b)  By  this  method  the  mixture  of  hydrogen,  marsh- gas,  and 
nitrogen  is  mixed  with  air  or  oxygen  sufficient  for  the  complete 
combustion  of  both  the  combustible  gases,  and  the  mixture  ex- 
ploded. The  contraction  is  then  measured,  after  which  the  carbon 
dioxide  is  absorbed  and  the  volume  again  measured.  From  these 
data  the  volumes  of  the  hydrogen  and  marsh-gas  can  be  calculated. 
As  already  explained,  p.  406,  the  contraction  due  to  the  combustion 
of  marsh-gas  is  twice  the  volume  of  carbon  dioxide  ;  if,  therefore, 
the  volume  of  carbon  dioxide  is  ascertained  (by  absorption  with 
potash),  and  twice  this  volume  be  deducted  from  the  contraction 
on  explosion,  the  product  will  represent  the  contraction  due  to  the 
combustion  of  the  hydrogen. 

Let  C  =  contraction  on  explosion,  and  C'  =  volume  of  CO2  pro- 
duced (i.e.  contraction  on  absorption  with  potash) ; 

Then  C  —  2C'  -  contraction  due  to  the  hydrogen 
and  |(C  -  2C)  =  volume  of  hydrogen 


410  Gas  Analysis. 

Again,  since  the  volume  of  carbon  dioxide  produced  is  the  same  as 
the  volume  of  marsh-gas  burnt — 

C  =  volume  of  marsh-gas 

EXAMPLE. — A  portion  of  the  mixture  of  hydrogen,  marsh-gas, 
and  nitrogen  employed  in  the  previous  example  (being  the  residual 
gas  after  the  removal  of  the  absorbable  constituents  from  a  sample 
of  coal-gas)  was  measured  in  the  burette,  and  an  excess  of  air 
added. 

Volume  of  gas  taken  =  14*2  c.c. 
volume  of  gas  +  air  =  97*6    „ 
after  explosion,  volume  =  74*2    „ 
therefore  contraction,  C,  =  97^6  —  74-2  =  23*4. 
After  absorption  by  KHO,  volume  =  68'2  c.c. 

therefore  C'  -  74*2  —  68*2  =  6 
Hence  volume  of  H  in  14-2  c.c.  of  the  gas  =  f  (23-4  —  12)  =  7*6  c  c. 

and  volume  of  CH4  in  14*2  c.c.  of  the  gas  =  6*O  ,. 
and  volume  of  N  in  14*2  c.c.  of  the  gas=  14^2 — (7§6-f  6-o)  =  O'6  „ 
Calculating  the  percentage  as  in  the  previous  example — 

7*6  x  86'6 

—  =  47*0  per  cent,  of  hydrogen 
14*2 

,  6-0  x  86'6 
and  ; =  36'5  per  cent,  of  marsh-gas 

0-6  x  86-6 

=  3'6  per  cent,  of  nitrogen 

14*2 

In  cases  where  the  gas  under  analysis  contains  a  relatively  large 
proportion  of  nitrogen  ;  as,  for  example,  in  the  case  of  producer 
gas,  or  blast-furnace  gases,  the  addition  of  air  would  dilute  the  gas 
to  such  an  extent  as  to  render  it  non-combustible.  Under  these 
circumstances,  therefore,  either  oxygen  must  be  substituted  for  air, 
or  else,  after  sufficient  air  has  been  added  to  furnish  the  requisite 
amount  of  oxygen,  a  few  cubic  centimetres  of  electrolytic  gas  may 
be  added  to  the  mixture.  The  electrolytic  gas  for  this  purpose 
may  be  generated  in  the  pipette  described  on  p.  404,  and  then 
transferred  to  the  mercury  explosion-pipette. 

The  Nitrometer. 

Many  simple  processes  of  gas  estimations  (such  as  when  it  is 
only  desired  to  determine  one  gas  by  absorption)  are  conveniently 
and  quickly  performed  by  means  of  the  Lunge  nitrometer.*  This 

*  So  called  because  originally  designed  by  Lunge,  for  the  estimation  of 
nitrogen  oxides  in  "  nitrous  vitriol." 


The  Nitrometer,  or  Gas-volumeter. 


411 


consists  of  a  calibrated  measuring-tube,  m  (Fig.  88),  connected  by 
means  of  stout  rubber  tubing  to  the  pressure-tube  p.  By  means 
of  the  two-way  tap  upon  the  measuring-tube,  communication  can 
be  established  at  will  either  with  the  bent  capil- 
lary tube  /  or  the  little  reservoir  n.  Mercury 
is  usually  employed  as  the  confining  liquid. 

To  introduce  the  gas,  the  pressure-tube  is 
raised  until  the  measuring-tube  is  completely 
rilled  with  mercury.  The  capillary  tube  /  is 
connected  to  the  supply  of  the  gas,  and  the 
two-way  tap  turned  so  as  to  open  communication 
between  the  measuring-tube  and  the  capillary, 
as  seen  in  the  figure,  and  the  gas  drawn  over 
by  lowering  the  pressure-tube.  By  turning  the 
tap  one  quarter  of  a  revolution,  communication 
with  both  the  exits  is  cut  off. 

The  absorbing  reagent  is  poured  into  the 
reservoir  »,  and  introduced  into  the  measuring- 
tube  by  first  slightly  lowering  the  pressure-tube, 
and  then  gently  turning  the  tap  so  as  to  open 
communication  between  m  and  ;/.* 

Besides  simple  absorption  operations,  this 
apparatus  is  suitable  for  the  estimation  of  the 
volume  of  gas  which  is  evolved  in  certain  defi- 
nite chemical  reactions.  These  processes  are 
sometimes  spoken  of  as  gas-volumetric  analysis. 
The  following  are  typical  examples  : — 

(i)  The  estimation  of  nitrates,  either  in 
commercial  nitre,  or  in  the  residue  obtained  on 
evaporating  water  for  the  determination  of  the  nitrates  and  nitrites. 

This  process  depends  upon  the  fact  that  when  a  nitrate  is 
decomposed  by  strong  sulphuric  acid  in  the  presence  of  mercury, 
this  metal  is  acted  upon  by  the  liberated  nitric  acid  with  the 
evolution  of  nitric  oxide.  The  nitric  oxide  is  therefore  the  measure 
of  the  nitric  acid  or  the  nitrate  present.  One  cubic  centimetre  NO 
measured  at  N.T.P.  represents  0-00452  gram  of  KNO3  or  0-0038 
gram  of  NaNO3. 

About  o*i  gram  of  the  nitrate  is  placed  in  the  reservoir  «,  and 
dissolved  in  2  or  3  c.c.  of  water.  This  solution  is  carefully  drawn 

*  As  a  safeguard,  lest  by  accident  the  tap  should  be  turned  the  wrong  way, 
a  rubber  connector  should  be  left  upon  the  end  of  /,  and  closed  with  a  pinch- 
cock. 


FIG. 


412  Gas  Analysis. 

into  the  measuring-tube  (previously  filled  with  mercury),  and  the 
reservoir  rinsed  with  I  c.c.  of  water,  which  is  also  allowed  to  pass 
into  the  tube.  (With  a  little  experience  these  operations  will  be 
performed  without  the  admission  of  any  air ;  but  should  air  be 
drawn  in,  it  may  be  expelled  by  raising  the  pressure-tube  slightly, 
and  cautiously  opening  the  tap.)  The  reservoir  is  again  rinsed  by 
the  introduction  of  5  or  6  c.c.  of  strong  sulphuric  acid,  which  is 
followed  by  the  addition  of  about  10  c.c.  more  acid.  The  measuring- 
tube  is  then  released  from  its  clamp,  and  the  contents  carefully 
shaken  by  inclining  the  tube,  and  then  quickly  (almost  with  a  jerk) 
bringing  it  into  the  vertical  position.  In  fifteen  to  twenty  minutes 
the  process  is  complete.  The  tube  is  replaced  in  its  clamp  and 
allowed  to  cool.  Before  reading  the  volume  of  gas,  in  order  to 
compensate  for  the  column  of  acid  in  the  tube,  a  similar  volume 
of  acid  of  the  same  dilution  is  poured  into  the  pressure-tube. 

(2)  By  means  of  a  short  rubber  tube,  the  tube  /  may  be  attached 
to  a  small  flask  fitted  with  a  caoutchouc  stopper  carrying  a  short 
glass  tube.  If  in  the  flask  a  chemical  reaction  resulting  in  the 
evolution  of  gas  at  the  ordinary  temperature  is  carried  out,  the 
volume  of  the  evolved  gas  can  be  measured. 

With  such  an  arrangement,  the  estimation  of  carbon  dioxide 
in  carbonates  may  be  made  ;  also  the  indirect  estimation  of 
manganese  dioxide,  by  measuring  the  carbon  dioxide  evolved  by 
the  action  of  the  dioxide  upon  oxalic  acid  in  presence  of  sulphuric 
acid,  according  to  the  equation — 

MnO2  +  H2SO4  +  C2H2O4  =  MnSO4  +  2H2O  +  2CO2 
Similarly,  by  the  action  of  hydrogen  peroxide  upon  manganese 
dioxide,  potassium  permanganate,  or  bleaching  powder,  the  oxygen 
evolved  is  a  measure  of  the  available  oxygen  in  these  compounds  ; 
half  the  oxygen  given  off  in  each  case  being  derived  from  the 
compound,  while  the  other  half  is  from  the  hydrogen  peroxide,  as 
seen  by  the  following  equations  : — 

MnO2  +  H2O2  =  MnO  +  H2O  +  O,  * 
2KMnO4  +  3H2SO4  +  5H202  =  K2SO4  +  2MnSO4-f  8H2O  +  5O2f 

Ca(OCl)Cl  +  H2O2  =  CaCl2  -f  H2O  +  O2 

A  weighed  quantity  of  the  substance  to  be  tested  is  introduced 
into  the  flask.  In  the  case  of  manganese  dioxide,  or  potassium 
permanganate,  dilute  sulphuric  acid  is  added,  but  with  bleaching 

*  In  the  presence  of  sulphuric  acid. 

f  The  importance  of  ensuring  the  presence  of  sufficient  sulphuric  acid  in 
this  and  the  preceding  reaction  will  be  seen  from  the  footnote  on  p.  54. 


The  Nitrometer^  or  Gas-volumeter.  413 

powder  a  small  quantity  of  water  only.  The  reagent  (in  the  above 
three  cases,  the  hydrogen  peroxide)  is  placed  in  a  test-tube,  and 
deposited  within  the  flask  without  allowing  any  of  it  to  come  in 
contact  with  the  materials  already  present.  The  rubber  cork  is 
inserted  in  the  mouth  of  the  flask,  and  the  apparatus  connected  to 
the  nitrometer. 

To  ensure  that  the  air  within  the  flask  is  under  atmospheric 
pressure,  the  two-way  tap  is  turned  so  as  to  open  communication 
between  the  flask  and  the  measuring-tube,  and  the  pressure-tube 
adjusted  so  that  the  mercury  stands  at  the  same  level  in  both 
tubes.  The  tap  is  then  turned  so  as  to  connect  the  reservoir  n 
with  the  measuring-tube,  and  the  air  from  the  latter  entirely 
expelled,  and  the  tube  filled  with  mercury  by  raising  the  pressure- 
tube.  The  tap  is  then  closed  and  the  pressure-tube  slightly 
lowered.  Communication  with  the  flask  is  again  established,  and 
the  contents  of  the  little  tube  tipped  out  into  the  flask  so  as  to 
bring  about  the  desired  reaction.  As  the  gas  is  evolved,  the 
pressure-tube  is  gradually  lowered  so  as  to  avoid  the  creation  of 
any  unnecessary  pressure  in  the  apparatus.  When  the  action  is 
completed,  the  mercury  is  brought  to  the  same  level,  and  the 
apparatus  allowed  to  stand  for  half  an  hour  for  the  gas  to  assume 
the  atmospheric  temperature,  when  the  volume  is  read  off.  The 
atmospheric  temperature  and  pressure  are  noted,  and  the  usual 
corrections  made  (see  p.  387). 


2  E 


PART  III. 

SECTION    I. 

ESTIMATION  OF   CARBON,  HYDROGEN,   NITROGEN,   CHLORINE, 
SULPHUR,  AND  PHOSPHORUS  IN  ORGANIC  COMPOUNDS. 

I.  Carbon  and  Hydrogen  by  Combustion. 

Epitome  of  Process. — These  two  elements  are  estimated 
simultaneously  by  one  and  the  same  series  of  operations,  whereby 
the  organic  compound  is  burnt,  and  the  carbon  and  hydrogen 
oxidised  to  carbon  dioxide  and  water  respectively.  The  oxidation, 
or  combustion,  of  the  organic  compound  is  accomplished  by  heating 
it  in  a  combustion-tube  with  copper  oxide,  in  a  gentle  stream  of  air 
or  oxygen.  The  products  of  the  combustion  are  thus  expelled  from 
the  tube,  and  are  made  to  pass  first  through  a  weighed  calcium 
chloride  tube,  and  then  through  weighed  potash  bulbs  (or  soda-lime 
tubes).  The  increase  in  weight  of  the  calcium  chloride  tube  is  due 
to  the  water,  one-ninth  of  which  represents  the  hydrogen  in  the 
organic  compound  ;  while  the  gain  in  weight  suffered  by  the  potash 
bulbs  is  caused  by  the  absorbed  carbon  dioxide,  three-elevenths  of 
which  is  the  weight  of  the  carbon  in  the  compound. 

Fitting  tip  the  Apparatus,  (i)  The  Combustion-tube. — 
A  piece  of  combustion-tube,  about  8  cms.  longer  than  the  com- 
bustion furnace  (which  itself  will  be  70  to  80  cms.  long*),  and  about 
1 5  mm.  bore,  is  rounded  at  the  edges  by  just  fusing  the  ends  in  the 
blowpipe  flame.  The  tube  is  then  thoroughly  cleaned  inside.  For 
this  purpose  one  end  is  closed  with  a  cork,  and  a  few  cubic  centi- 
metres of  strong  sulphuric  acid  poured  in  at  the  other  end.  This  is 
made  to  flow  over  the  whole  interior  surface  by  tipping  and  rotating 
the  tube.  The  acid  is  then  poured  away,  and  the  tube  thoroughly 
rinsed  out  with  water,  the  cork  being  removed  and  a  stream  of  water 
run  through  the  tube.  The  tube  is  allowed  to  drain,  and  finally 
dried  by  gently  warming  it  and  blowing  air  through  it  from  the 
bellows. 

*  For  rough  measures,  2 '5  cms.  —  i  inch. 


Carbon  and  Hydrogen  Determination.          415 


Into  one  end  of  the  tube  a  roll  of  copper  gauze,  c,  Fig.  89,  is 
thrust  to  a  distance  of  about  20  cms.  This  is  made  by  rolling  a 
strip  of  fine  gauze,  2*5  cms. 
wide,  into  a  cylinder  or  plug, 
which  will  fit  moderately 
tightly  into  the  tube,  so  as  to 
retain  its  place.  The  tube  is 
then  filled  up  to  within  12 
cms.  of  the  other  end  with 
granular  copper  oxide  ;  *  and 
a  second  roll  of  copper  gauze, 
c,  the  same  length  as  the 
other,  is  pushed  into  the  tube 
so  as  to  confine  the  copper 
oxide. 

The  substance  under  ana- 
lysis is  contained  in  a  platinum 
or  porcelain  boat,  £,  which  is 
pushed  into  the  tube  nearly 
close  up  to  the  copper-gauze 
roll. 

Behind  the  boat  is  intro- 
duced a  loose-fitting  plug,£>, 
the  object  of  which  is  to  con- 
centrate the  current  of  air  or 
oxygen  at  this  point,  in  order 
to  prevent  the  backward  diffu- 
sion of  vapours  or  products 
of  combustion  of  the  sub- 
stance in  the  boat.  This  plug 
may  be  either  a  short  roll  of 
copper  gauze,  having  a  little 
wire  loop  by  means  of  which 
it  can  easily  be  withdrawn  ; 
or,  better,  a  short  piece  of 
hard  glass  tube  (which  will 
just  pass  freely  into  the  com- 
bustion-tube), closed  at  one 
end,  and  narrowed,  but  left 

*  Copper  oxide  made  by  heating  the  nitrate  should  not  be  used,  as  it  is 
liable  to  evolve  oxides  of  nitrogen  when  heated.  Granular  copper  oxide  is 
obtained  by  direct  oxidation  of  the  metal. 


416  Ultimate  Organic  Analysis. 

open,  at  the  other  end.  This  can  readily  be  withdrawn  by  means 
of  a  piece  of  wire,  bent  into  a  small  hook  at  the  end. 

The  combustion-tube  is  closed  at  this  end  with  a  rubber  stopper 
carrying  a  short  piece  of  narrow  glass  tube,  which  projects  about 
3  mm.  into  the  combustion-tube,  the  other  end  being  connected  to 
the  apparatus  for  purifying  and  drying  the  air  and  oxygen.  The 
space,  S,  at  the  other  end  of  the  tube  is  for  the  reception  of  a  roll 
of  either  copper  or  silver  gauze,  about  6  cms.  long,  and  sufficiently 
loose  to  admit  of  easy  removal.  If  the  organic  compound  to  be 
analysed  contains  nitrogen,  either  the  copper  or  silver  roll  will 
serve  to  decompose  the  nitrogen  peroxide,*  which,  if  allowed  to  pass 
from  the  tube  undecomposed,  would  be  absorbed  by  the  potash,  and 
thereby  vitiate  the  carbon  determination.  When,  however,  the 
compound  contains  a  halogen,  the  silver  roll  must  be  employed. 

Before  use,  such  a  copper  gauze  roll  should  be  first  superficially 
oxidised,  by  heating  for  a  few  moments  in  a  flame,  and  then  heated 
in  a  stream  of  hydrogen.  A  number  of  rolls  may  be  prepared  at 
one  operation,  and  may  be  kept  bright  by  being  corked  in  a  tube. 

(2)  The  Combustion  Furnace. — When  the  combustion-tube   is 
fitted  up,  it  is  laid  in  the  furnace,!  and  should  project  about  3^  cms. 
at  each  end.     A  shallow  trough  or  tray  of  sheet  iron  runs  the  entire 
length  of  the  furnace,  and  in  this  a  bed  for  the  tube  is  made  by  first 
laying  in  the  trough  a  strip  of  asbestos  cloth,  wide  enough  to  pro- 
ject over  the  sides  of  the  iron  trough,  so  that  the  tube  shall  rest 
uniformly  upon  the  asbestos  without  coming  in  contact  with  the 
iron. 

(3)  The  Calcium  Chloride  Tufa.— This  is  filled  and  prepared 
precisely  as  described  on  p.  203. J     When  not  in  use,  the  ends  of 
the  tube  are  closed  by  short  rubber  tubes  plugged  with  pieces  of 
glass  rod.     These  caps  are  always  removed  while  the  tube  is  being 
weighed,  but  immediately  replaced  until  the  apparatus  is  required 
for  the  combustion. 

The  calcium  chloride  tube  is  attached  to  the  combustion-tube 
by  means  of  a  rubber  stopper,  and  is  conveniently  supported  by  a 
retort  stand  and  ring,  as  shown  in  the  figure. 

*  4Cu  +  2N02  =  4CuO  +  N2. 

f  Erlenmeyer's  gas  furnace,  which  consists  practically  of  a  row  of  Bunscn 
burners,  is  one  of  the  best  forms. 

t  As  samples  of  calcium  chloride  are  sometimes  slightly  alkaline,  it  is 
desirable,  when  the  tube  is  freshly  filled,  to  pass  a  stream  of  dry  carbon  dioxide 
through  it  for  a  few  minutes,  and  then  to  displace  the  carbon  dioxide  by  a 
current  of  dry  air,  which  should  be  passed  through  the  tube  for  about  half  an 
hour. 


Carbon  and  Hydrogen  Determination.          417 


(4)  The   Potash   Bulbs. —These  may  be  either   the  ordinary 
Liebig's  bulbs,  or  the 

more  modern  Geissler's 
apparatus,  shown  at  B, 
Fig.  89.  They  are 
charged  with  a  strong 
solution  of  caustic 
potash  by  dipping  the 
tube  projecting  from 
the  larger  bulb  into  a 
small  beaker  or  dish 
containing  the  potash, 
and  sucking  the  liquid 
into  the  apparatus  by 
means  of  a  piece  of 
rubber  tube  upon  the 
opposite  end.  A  little 
guard  tube,  G,  filled 
with  fragments  of  solid 
caustic  potash,  is  at- 
tached to  the  bulbs  in 
order  to  prevent  the 
loss  of  water-vapour, 
which  would  otherwise 
result  from  the  passage 
through  the  solution 
of  the  dry  gases  passing 
from  the  combustion- 
tube.  Just  as  with  the 
calcium  chloride  tube, 
when  not  in  actual  use 
the  ends  of  the  tubes 
are  closed  with  rubber 
caps,  which  are  re- 
moved while  the  appa- 
ratus is  being  weighed. 

(5)  The  Apparatus 
for  Purifying  the  Air 
and     Oxygen.  —  This 
apparatus      must     be 
fitted  up  in  duplicate, 

so  that  either  air  or  oxygen  can  be  admitted  into  the  combustion- 


41 8  Ultimate  Organic  Analysis. 

tube  in  a  purified  condition  without  the  loss  of  time  which  would 
result  if  one  gas  were  made  to  follow  the  other  through  the  whole 
apparatus.  Fig.  90  shows  a  convenient  arrangement  for  the  purpose. 

Each  gas  (contained  in  gas-holders,  or  compressed  in  cylinders) 
is  passed  first  through  a  two-necked  bottle,  B,  B',  nearly  filled  with 
pumice  moistened  with  sulphuric  acid.  They  then  pass  through 
the  eprouvettes  E,  E',  which  are  filled  with  fragments  of  solid 
caustic  potash.  This  part  of  the  apparatus,  when  once  charged,  may 
remain  undisturbed  for  several  months,  and  may  conveniently  be 
placed  upon  a  shelf  out  of  the  way.  The  gases  may  be  conducted 
to  it  and  led  away  from  it  by  means  of  fine  "  compo  "  pipe  3  mm. 
in  the  bore,  which  may  be  carried  down  the  wall  to  the  side  of  the 
draught  chamber,  in  which  the  combustion  furnace  is  placed.  At 
this  point  they  are  joined  to  two  limbs  of  a  bridle-tube,  T,  by  means 
of  short  rubber  connections,  each  with  its  screw-clamp  or  stop-cock  ; 
these  joints  should  all  be  securely  wired.  The  third  limb  of  the 
bridle-tube  is  attached  to  a  U-tube  filled  with  soda-lime,  and  this  in 
its  turn  is  connected  to  a  set  of  Liebig's  bulbs  containing  strong 
sulphuric  acid.  By  manipulating  the  taps,  either  oxygen  or  air  can 
be  admitted  into  these  two  tubes,  and  the  rate  at  which  the  gas 
passes  is  seen  by  the  bubbles  travelling  through  the  bulbs  containing 
the  sulphuric  acid. 

Testing  the  Apparatus.— When  the  apparatus  is  fitted  up, 
it  is  necessary,  before  proceeding  to  make  an  actual  combustion, 
to  ensure  the  complete  removal  from  the  combustion-tube  of  every 
trace  of  moisture  and  of  combustible  matter.  The  calcium  chloride 
tube  and  the  potash  bulbs  are  for  this  purpose  removed,  and  the 
roll  of  copper  gauze  is  introduced  into  the  space  S  (Fig.  89).*  The 
boat  and  the  glass  plug  p  are  also  removed. 

The  combustion-tube  is  then  heated  throughout  its  entire  length 
to  a  low  red  heat  for  about  20  minutes,  while  a  slow  stream  of  the 
dried  and  purified  air  is  passed  through  it.  The  air-supply  should 
be  then  stopped,  and  the  air  contained  in  the  oxygen  purifying 
apparatus  made  to  pass  through  the  tube  by  sweeping  it  out  with 
a  stream  of  oxygen.  This  may  be  passed  rather  more  rapidly  than 
at  first,  so  that  in  about  15  minutes  oxygen  will  be  issuing  from  the 
extreme  end  of  the  tube. 

*  It  is  advisable  to  place  this  roll  in  the  tube  at  this  stage,  even  although 
the  combustions,  which  are  to  be  first  made,  are  of  substances  which  do  not 
contain  nitrogen;  so  that  afterwards,  when  the  copper  maybe  required,  its 
freedom  from  organic  matter  will  be  assured.  Moreover,  in  case  the  ordinary 
copper  oxide  (prepared  from  the  nitrate)  is  the  only  preparation  at  hand,  the 
presence  of  this  copper  roll  will  serve  to  decompose  the  oxides  of  nitrogen  which 
are  derived  from  this  source. 


Carbon  and  Hydrogen  Determination.          419 

While  this  is  going  on,  the  calcium  chloride  tube  and  the  potash 
bulbs  are  weighed,  the  caps  being  removed  for  the  short  time 
occupied  by  the  weighing. 

The  passage  of  the  oxygen  is  then  stopped,  and  the  weighed 
tubes  attached,  the  ends  of  the  glass  tubes  of  the  two  pieces 
which  are  connected  by  the  rubber  union,  being  pushed  up  within 
the  rubber  so  as  to  touch  each  other.* 

A  slow  stream  of  air  is  now  passed  through  the  apparatus,  the 
rate  being  that  which  would  be  adopted  during  an  actual  combustion, 
namely,  about  two  bubbles  per  second,  as  they  pass  through  the 
bulbs.  This  is  continued  for  about  15  minutes,  after  which  the 
stream  of  air  is  replaced  by  a  stream  of  oxygen  at  about  the  same 
rate  for  another  15  minutes.  The  potash  bulbs  and  calcium 
chloride  tube  are  then  carefully  disconnected,  and  their  ends  im- 
mediately closed  with  their  caps.  After  a  short  interval,  during 
which  they  are  allowed  to  assume  the  ordinary  temperature,  they 
are  weighed  (as  usual,  without  the  caps),  and  if  the  operation  has 
been  successfully  carried  out,  their  weight  should  have  undergone 
no  increase.  If,  on  the  other  hand,  either  piece  of  apparatus  has 
gained  more  than  3  or  4  milligrams,  it  shows  that  the  first  part  of  the 
operation  has  not  been  complete,  and  that  either  moisture  or  carbon 
dioxide  is  still  being  evolved.  In  this  case  the  passage  of  air  through 
the  still  heated  tube  must  be  continued  for  a  short  time,  and  another 
"  blank  "  experiment  made  by  replacing  the  calcium  chloride  tube 
and  potash  bulbs,  and  allowing  first  air,  and  then  oxygen,  to  pass 
through  the  entire  apparatus  for  15  minutes  each.  The  absorption 
apparatus  is  then  detached  and  reweighed,  when,  unless  something 
is  seriously  wrong,  the  weights  will  be  found  to  have  undergone  no 
change. 

The  apparatus  is  now  in  a  condition  to  be  used  for  a  series  of 
"  combustions,"  and  need  not  be  further  tested  until  it  becomes 
necessary  to  replace  the  combustion-tube.  If  the  asbestos  bed  has 
been  properly  arranged,  so  that  the  tube  rests  uniformly  all  along 
it,  and  if  care  is  exercised  in  heating  the  tube,  so  as  not  to  heat  too 

*  It  is  of  the  greatest  importance  that  the  stopper  or  cork  through  which  the 
calcium  chloride  tube  passes  should  not  become  over-heated  during  a  ' '  com- 
bustion," and  yet  at  the  same  time  the  tube  must  be  sufficiently  warm  right  up 
to  the  cork  to  prevent  moisture  from  condensing  before  it  passes  into  the  bulb. 
If  the  combustion-tube  projects  4  cms.  beyond  the  furnace,  as  directed  on  p.  416, 
the  cork  is  practically  quite  safe,  but  the  additional  precaution  may  be  adopted 
of  placing  a  small  screen,  cut  out  of  asbestos  cardboard,  over  the  tube  just 
outside  the  furnace.  This  screen,  however,  should  be  removed  towards  the 
end  of  the  process  to  make  sure  that  no  moisture  has  deposited  before  passing 
into  the  absorption-tube. 


42O  Ultimate  Organic  Analysis. 

suddenly  at  first,  and  not  to  raise  the  temperature  unnecessarily 
high,  whereby  the  glass  is  softened,  a  combustion-tube  made  of 
the  best  hard  glass  will  stand  being  used  over  and  over  again  for 
a  great  many  analyses. 

Typical  Examples  of  Combustions. 

(a)  Combustion  of  Non-volatile  Solids,— One  of  the  most 
suitable  substances  to  employ  for  a  first  experiment  is  cane  sugar, 
as  it  is  readily  obtained  in  a  state  of  purity,  and  the  percentage  of 
hydrogen  it  contains  is  comparatively  high.  The  pure  crystallised 
compound  is  finely  powdered,  and  dried  in  a  steam-oven. 

A  porcelain  or  platinum  boat  (previously  proved  to  be  of  such 
a  size  that  it  will  pass  easily  into  the  combustion-tube)  is  supported 
on  a  pipeclay  triangle,  and  heated  to  low  redness  by  means  of  a 
Bunsen  flame  ;  if  a  porcelain  boat  is  employed,  it  must  be  heated 
cautiously. 

It  is  then  placed  to  cool  in  the  desiccator,  and  when  cold,  it  is 
weighed.  It  must  be  handled  as  little  as  possible  by  the  fingers, 
and  not  at  all  if  the  hands  of  the  operator  are  at  all  moist. 

From  o'2  to  0^3  gram  of  the  dry  sugar  is  weighed  into  the  boat, 
which  is  then  pushed  into  the  combustion-tube,  almost  close  up  to 
the  roll  of  copper  gauze.  The  clean  dry  glass  plug  is  next  intro- 
duced, and  the  tube  is  then  connected  to  the  apparatus  for  supplying 
the  stream  of  dry  and  pure  air. 

The  weighed  calcium  chloride  tube  and  potash  bulbs  are  then 
attached,  the  projecting  end  of  the  combustion-tube  being  firmly 
held  with  one  hand  while  the  calcium  chloride  tube  is  being  fitted 
on,  so  as  not  to  disturb  the  boat  and  its  contents. 

The  portion  of  the  tube  from  S  (Fig.  89)  to  the  copper  gauze 
roll  c,  that  is,  the  part  containing  the  whole  of  the  copper  oxide,  is 
first  heated  to  redness,  after  which  a  slow  stream  of  air  (not  faster 
than  two  bubbles  per  second)  is  passed  through.  The  heating  is 
now  gradually  extended  towards  the  boat  by  turning  on  one  burner 
at  a  time — at  first  with  a  small  flame,  and  then  increasing  it. 

The  success  of  the  operation  hinges  upon  this  part  of  the 
process,  therefore  it  must  be  carefully  watched,  lest  the  heat  be 
applied  too  rapidly. 

The  combustion  of  sugar  is  usually  completely  effected  with  a 
stream  of  air,  without  the  admission  of  pure  oxygen  ;  but  if  any 
traces  of  a  black  residue  are  seen  to  persist  in  the  boat  towards  the 
end  of  the  operation,  a  stream  of  oxygen  may  be  substituted  for 


Typical  Examples  of  Combustions.  421 

the  air  fora  few  minutes,  after  which  the  oxygen  is  stopped,  and  the 
air  again  passed  through  for  about  10  or  15  minutes. 

In  this  way  the  whole  of  the  water-vapour  and  carbon  dioxide 
is  swept  out  of  the  tube  into  the  absorption  apparatus  ;  and  the 
portions  of  copper  oxide  which  during  the  combustion  had  become 
reduced  are  re- oxidised. 

The  potash  bulbs  and  calcium  chloride  tube  are  then  removed, 
their  ends  capped,  and,  after  being  allowed  to  assume  the  ordinary 
temperature,  they  are  weighed. 

H2O  :  H2  =  9  :  i 
18       2 

f       weight  of  H9O 
therefore  —  -«•  =  weight  of  hydrogen 

and  CO2  :  C  =  1 1  :  3 
44       12 

f       weight  of  CO9  x  3  .  ,       r 

therefore *  =  weight  of  carbon 

From  the  numbers  thus  obtained,  the  percentage  of  hydrogen 
and  carbon  are  calculated.  The  oxygen  is  estimated  by  difference.* 

The  theoretical  composition  of  sugar,  calculated  from  the 
formula  C12H22On,  is  hydrogen,  6*43  ;  carbon,  42*10  ;  oxygen,  51*47. 

The  experimental  errors  of  a  carbon  and  hydrogen  determination 
are  usually  in  the  direction  of  a  deficit  of  the  former,  and  a  slight 
excess  of  the  latter,  element.  If,  in  the  analysis  of  sugar,  the  carbon 
falls  more  than  0-3  below  the  theoretical  value,  or  the  hydrogen 
more  than  o-i  above  it,  a  duplicate  experiment  should  be  made. 

(£)  Combustion  of  Volatile  Solids. — For  practice  in  the 
combustion  of  a  volatile  solid,  benzoic  acid  may  be  employed. 

Benzoicacid,  CGH5.COOH  :  carbon,  68*85 2  per  cent. ;  hydrogen, 
4*918  per  cent. 

From  o!4  to  o*  5  gram  of  the  compound  is  weighed  out  into  the 
boat,  but  before  the  latter  is  introduced  into  the  combustion-tube, 
the  weighed  calcium  chloride  tube  and  potash  bulbs  are  attached, 
and  the  tube  is  heated  to  redness  to  within  about  2|  cms.  of  the 
copper  roll  c  (Fig.  89).  The  boat  is  then  carefully  pushed  in,  the 
glass  plug  quickly  inserted,  and  connection  made  with  the  gas- 
drying  apparatus.  A  slow  stream  of  air  is  admitted,  and  then  the 
heating  very  gradually  and  cautiously  extended  in  the  direction  ol 

*  In  cases  where  the  substance  under  analysis  leaves  an  incombustible 
residue  or  ash,  the  boat  is  withdrawn  after  the  tube  has  cooled,  and  placed  in 
a  desiccator  until  cold,  when  it  is  weighed.  Deducting  the  weight  of  the  boat 
gives  the  weight  of  the  ash. 


422  Ultimate  Organic  Analysis, 

the  boat.  The  more  volatile  the  substance  is,  the  more  gradually 
must  the  heating  be  extended  towards  the  boat  containing  it. 

The  remainder  of  the  process  is  conducted  as  in  the  former 
example,  the  stream  of  air  being  replaced  by  pure  oxygen  towards 
the  end  of  the  operation. 

(c]  Combustion  of  Liquid  Compounds.— If  the  liquid  is 
not  appreciably  volatile  at  ordinary  temperatures,  and  at  the  same 
time  does  not  absorb  atmospheric  moisture,  it  may  be  weighed  out 
into  the  boat  and  treated  as  in  the  above  example. 

Moderately  volatile  liquids  are  enclosed  in  small  glass  bulbs, 
made  by  drawing  out  a  piece  of  narrow  glass  tube,  as  shown  at  A 
(Fig.  91),  and  then  blowing  the  small  tube  into  a  bulb,  B.  The 

bulb  must  be  of  such  a  size  that 
it  will  lie  in  the  boat,  and  pass 
into  the  combustion-tube  in  this 
position. 

As    an     example,     benzene, 

A  CGH6,  may  be  employed.     One 

FlG-  9J.  of  the  little  bulbs,  with  a  capil- 

lary stem  about  3  cms.  long,  is  weighed.  It  is  then  gently  warmed, 
and  the  stem  dipped  into  the  liquid  in  a  small  beaker.  When  about 
025  gram  has  entered  the  bulb,  the  tip  of  the  stem  is  sealed  by 
bringing  it  into  a  small  gas-flame.  The  bulb  is  then  weighed. 

As  in  the  case  of  the  volatile  solid,  the  absorption  apparatus  is 
attached,  and  the  combustion-tube  is  raised  to  a  red  heat  to  within 
about  4  cms.  of  the  copper  coil  c  (Fig.  89)  before  the  boat  containing 
the  bulb  is  introduced. 

A  very  slight  scratch  with  a  fine-cut  file  is  made  on  the  capillary 
stem  of  the  bulb,  near  to  the  closed  end,*  and  the  bulb  is  laid  in 
the  boat  in  such  a  position  that  the  fine  point  projects  over  the  end 
of  the  boat  beyond  the  file-mark.  In  this  way  the  liquid  is  pre- 
vented from  running  up  into  the  capillary  tube.  When  all  is  in 
readiness,  the  end  of  the  fine  tube  is  broken  off,  the  fragment  is 
deposited  in  the  boat  (in  order  that  any  traces  of  liquid  it  contains 
may  be  included),  and  the  latter  immediately  pushed  into  the 
combustion-tube,  the  stem  of  the  bulb  foremost.  The  plug  is 
quickly  introduced,  and  the  tube  connected  with  the  gas-drying 
apparatus.  The  process  is  then  continued  as  in  the  former  example. 

*  The  liquid  should  not  be  more  in  the  capillary  tube  than  can  be  avoided. 
It  may  be  driven  back  into  the  bulb  so  as  to  have  a  clear  place  near  the  end 
where  the  stem  is  afterwards  to  be  broken,  by  gently  warming  the  tube  with  a 
very  small  flame. 


Combustion  of  Volatile  Liquids.  423 

Liquids  which  are  highly  volatile  are  vaporised  outside  the 
combustion-tube,  and  the  vapour  slowly  passed  over  the  heated 
copper  oxide.  For  this  purpose  a  piece  of  moderately  narrow  glass 
tube  is  drawn  out  into  thick-walled  capillaries  in  two  places,  so  as 
to  leave  a  piece  of  the  original  tube  about  2^  cms.  (i  inch)  long. 
In  order  that  the  walls  of  the  capillary  tubes  shall  be  thick,  the 
glass  tube  is  heated  in  the  blowpipe  over  as  long  a  surface  as  the 
flame  will  allc;w,  until  the  walls  of  the  tube  have  softened  and 
thickened  as  much  as  possible  without  falling  together.  The  tube 
is  then  slowly  drawn  out  the  desired 
length,  about  8  cms.  (3  inches).  It 
is  then  bent  in  the  manner  shown 
in  Fig.  92.  The  tube  is  first 
weighed,  after  which  about  o'3  to  0-5 
gram  of  the  liquid  is  introduced, 
and  the  fine  drawn-out  ends  sealed 
in  a  flame.  The  tube  is  then  re- 
weighed.  The  liquid  must  not  be 
allowed  to  get  up  into  the  hori- 
zontal capillary  tube.  A  cork, 
previously  selected  to  fit  the  com- 
bustion-tube, and  dried  in  an  air- 
oven,  is  fitted  on  to  the  horizontal 
limb  of  the  little  tube,  and  the 

apparatus  is  cooled  by  being  immersed  in  ice-water,  or  in  a  freezing 
mixture  (depending  upon  the  volatility  of  the  liquid),  so  as  to  reduce 
the  tension  of  its  vapour  to  an  insignificant  pressure.  Thus,  for 
such  a  liquid  as  ether,  ice-water  may  be  used. 

The  weighed  absorption  apparatus  having  been  attached,  and 
the  combustion-tube  heated  to  redness,  the  extreme  end  of  the 
capillary  tube  projecting  through  the  cork  is  cut  off,  and  the  cork 
inserted.  The  glass  plug  is  not  introduced  in  this  case. 

The  temperature  of  the  little  tube  containing  the  liquid  is  then 
allowed  gradually  to  rise.  If  it  is  in  ice- water,  this  may  be  done  by 
gently  applying  heat  to  the  water  by  bringing  a  small  flame  under 
the  beaker,  while,  if  it  is  in  a  freezing  mixture,  the  gradual  addition 
of  ordinary  water  will  be  sufficient  to  raise  the  temperature. 

In  this  way  the  liquid  is  slowly  vaporised,  and  the  vapour 
caused  to  pass  over  the  heated  copper  oxide.  When  the  liquid  has 
entirely  volatilised,  a  fine  scratch  is  made  with  a  file  close  to  the 
sealed  end  of  the  apparatus,  and  the  rubber  tube  from  the  gas- 
purifying  apparatus  is  slipped  on.  The  end  is  then  snapped  off 


424  Ultimate  Organic  Analysis. 

inside  the  rubber,  and  a  gentle  stream  of  air  (afterwards  followed 
by  oxygen  if  the  organic  compound  is  rich  in  carbon)  is  passed 
through  the  tube  until  the  combustion  is  complete,  and  the  products 
have  been  entirely  expelled  into  the  absorption  apparatus. 

II.  Estimation  of  Nitrogen  in  Organic  Compounds. 
Nitrogen  in  organic  substances  may  be  estimated  either  by  burning 
the  compound  with  copper  oxide  in  a  combustion-tube  (very  much 
as  in  a  carbon  and  hydrogen  determination  already  described),  and 
collecting  and  measuring  the  nitrogen  that  is  evolved  ;  or  it  may 
be  indirectly  determined  by  converting  it  into  ammonia,  and  esti- 
mating the  ammonia  by  one  of  the  processes  described  in  former 
sections.  The  former  process  (Dumas)  is  applicable  to  all  nitro- 
genous organic  compounds,  while  the  latter  cannot  be  employed  in 
the  case  of  those  substances  in  which  the  nitrogen  is  present  in  a 
"  nitro  "  or  "  azo  "  group. 

i.  Nitrogen  by  Dumas'  Method. 

Epitome  of  Process. — The  organic  compound  *  is  mixed 
with  copper  oxide  in  a  combustion-tube  closed  at  one  end,  contain- 
ing at  the  closed  end  a  quantity  of  pure  sodium  bicarbonate. 
(This  furnishes  a  supply  of  carbon  dioxide  with  which  to  sweep 
out  first  the  air,  and  finally  the  products  of  the  combustion,  from 
the  tube.)  A  roll  of  bright  copper  gauze  is  introduced  at  the  open 
end,  to  prevent  the  escape  of  any  nitrogen  in  the  form  of  oxides  of 
nitrogen.  The  tube  is  closed  with  a  cork  and  delivery  tube,  and 
the  nitrogen  is  collected  in  a  measuring-tube  over  caustic  potash, 
whereby  the  carbon  dioxide  is  removed. 

A  piece  of  combustion-tube  (cleaned  in  the  manner  described 
on  p.  414)  is  drawn  off  and  sealed  at  one  end  before  the  blowpipe. 
The  tube  should  be  about  76  cms.  (30  inches  f)  long,  or  of  such  a 
length  that  it  can  be  heated  to  the  extreme  end  when  in  the  furnace. 

A  quantity  of  pure  dry  sodium  bicarbonate  %  is  introduced,  sc 
that  when  the  closed  end  is  gently  tapped  on  the  table,  the  salt 
shall  occupy  a  space  of  about  20  cms.  A  layer  of  about  12  cms.  of 
powdered  copper  oxide  is  introduced,  the  tube  being  conveniently 
supported  by  a  clamp  in  a  vertical  position.  About  0*5  gram  of  the 
organic  substance  is  then  carefully  introduced  into  the  tube  (the 

*  A  suitable  compound  to  employ  for  practice  in  the  process  is  urea, 
CO(NH2)2. 

f  For  rough  measures,  2*5  cms.  -  i  inch. 

\  Sodium  bicarbonate  of  commerce  sometimes  contains  appreciable  quanti- 
ties of  ammonia,  when  manufactured  by  the  ammonia-soda  process.  If  the 
sample  to  be  used  is  of  unknown  origin,  a  "blank"  experiment  should  first  be 
made,  in  which  sugar  or  some  other  pure  organic  compound  containing  no 
nitrogen  is  used  instead  of  urea. 


Nitrogen  by  Dumas'  Method.  425 

weighing-bottle  containing  the  compound  is  first  weighed,  and  after 
a  suitable  quantity  has  been  transferred  to  the  tube,  the  bottle  is 
reweighed),  and  by  means  of  a,  long  stout  clean  wire,  bent  like  a 
corkscrew  at  the  end,  this  is  thoroughly  stirred  into  the  copper 
oxide  below,  to  a  distance  of  about  one-half  the  depth  of  the  copper 
oxide.  A  little  more  copper  oxide  (about  2^  cms.)  is  added  (with- 
out withdrawing  the  wire),  and  the  top  layers  of  the  lower  stratum 
are  mixed  upwards  into  this.  Once  more  a  similar  quantity  of 
copper  oxide  is  added,  and  the  corkscrew  thoroughly  freed  from 
any  traces  of  the  organic  substance  by  being  twisted  up  through 
this. 

The  tube  is  then  rilled  to  within  about  12  cms.  of  the  top  with 
granular  copper  oxide,  and  lastly,  a  tolerably  tight-fitting  roll  of 
bright  copper  gauze,  7  cms.  long,  is  introduced.  The  tube  is  then 
removed  from  the  clamp,  and,  while  held  in  a  horizontal  position 
by  the  thumb  and  fore  finger  of  both  hands,  it  is  gently  tapped  upon 
the  table.  In  this  way  the  contents  are  made  to  settle  down,  so  as 
to  leave  a  narrow  channel  or  air-space  all  along  the  top.  The  tube 
is  then  laid  in  the  furnace. 

A  well-fitting  cork  with  a  bent  delivery  tube  is  inserted,  and 
that  portion  of  the  tube  containing  the  granular  copper  oxide  is 
heated  to  dull  redness. 

As  soon  as  the  tube  is  visibly  red  hot,  the  most  forward  portions 
of  the  sodium  carbonate  are  heated  so  as  to  sweep  out  the  air  from 
the  tube  by  means  of  the  carbon  dioxide  so  generated.  Not  more 
than  half  the  carbonate  should  be  thus  used  at  this  stage. 

As  soon  as  the  heating  of  the  carbonate  is  commenced,  the 
delivery  tube  is  attached,  by  means  of  a  rubber  connection,  to  the 
lower  branch  tube  of  a  SchifFs  burette,  as  shown  in  Fig.  93.  At 
the  bottom  of  the  burette  there  is  previously  introduced  a  small 
quantity  of  mercury,  sufficient  to  reach  about  12  mm.  (|  inch) 
above  the  junction  of  the  branch  tube,  and  a  strong  solution  of 
caustic  potash  (i  part  of  solid  KHO  in  2  parts  of  water)  is  poured 
into  the  movable  reservoir  R.  The  tap  /  is  then  opened,  and  the 
burette  filled  with  potash  solution  by  raising  the  reservoir,  after 
which  the  tap  is  again  closed,  and  the  reservoir  placed  in  about  the 
position  shown  in  the  figure.  So  long  as  air  is  still  being  expelled 
from  the  combustion-tube,  gas  will  collect  in  the  burette  ;  but  as  it 
becomes  swept  out  by  the  carbon  dioxide,  the  ascending  bubbles 
get  smaller  and  smaller,  until  at  last  they  are  practically  entirely 
absorbed  by  the  potash.  When  this  point  is  reached,  the  air  which 
has  collected  is  expelled  by  again  cautiously  raising  the  reservoir 


426 


Ultimate  Organic  Analysis. 


R 


and  opening  the  tap,  and  the  heating  of  the  combustion-tube  ex- 
tended first  to  the  roll  of  copper  gauze,  and  then  gradually  along  the 

tube  towards  the  mixture  of 
copper  oxide  and  the  organic 
substance. 

The  nitrogen  which  is 
evolved,  together  with  the 
carbon  dioxide  now  rilling  the 
air-space  of  the  tube,  pass  up 
into  the  burette,  and,  the  car- 
bon dioxide  being  absorbed 
by  the  potash,  the  nitrogen 
alone  collects.  When  the 
evolution  of  nitrogen  is  com- 
pleted, the  remainder  of  the 
sodium  bicarbonate  is  heated, 
whereby  a  fresh  supply  of 
carbon  dioxide  is  generated, 
which  drives  out  the  remain- 
der of  the  nitrogen  now  rilling 
the  tubes.  This  is  continued 
until  the  bubbles,  which  enter 
the  burette  through  the  mer- 
cury, are  absorbed  as  before 
by  the  potash. 

In  order  to  make  sure  that  the  carbon  dioxide  is  completely 
removed,  the  cup  c  is  filled  up  with  fresh  potash  solution,  which  is 
slowly  admitted  into  the  burette  by  cautiously  opening  the  tap. 
When  the  operation  is  completed,  the  delivery  tube  is  withdrawn 
from  the  rubber  connector,  which  is  then  closed  with  a  pinch-cock.* 
The  reservoir  is  raised  until  the  surface  of  the  liquid  it  contains  is 
level  with  that  in  the  burette,  in  which  position  the  apparatus  is 
left  for  about  a  quarter  of  an  hour,  to  ensure  the  perfect  absorption 
of  any  traces  of  carbon  dioxide  which  may  be  present.  The  levels 
are  then  exactly  adjusted,  and  the  volume  of  the  nitrogen  read  off 
upon  the  graduated  tube.f 

*  In  some  forms  of  Schiffs  burette,  the  branch  tube  is  sufficiently  wide  to 
admit  of  its  being  closed  with  a  small  cork  ;  but  in  this  case  it  is  less  easy  to 
secure  an  air-tight  connection  with  the  delivery  tube. 

t  Where  a  Schiff  s  burette  is  not  at  hand,  the  gas  may  be  collected  in  a 
graduated  glass  tube  filled  with  strong  potash  solution,  and  inverted  in  a  trough 
of  mercury.  The  point  at  which  the  air  in  the  combustion-tube  has  been  all 
swept  out  by  carbon  dioxide,  must  in  this  case  be  ascertained  by  placing  the 


Nitrogen.  427 

The  atmospheric  temperature  and  pressure  are  noted,  and  the 
volume  of  gas  reduced  to  N.T.P.  by  the  formula  on  p.  388.  As  the 
tension  of  aqueous  vapour  exerted  by  such  a  strong  solution  of 
potash  as  is  here  used  is  considerably  less  than  that  of  water  alone, 
it  is  usual  to  give  to  p  (see  p.  388)  the  value  of  half  the  tension 
of  vapour  of  water,  taken  from  the  table  in  the  Appendix. 

Since  I  c.c.  of  nitrogen,  under  normal  conditions,  weighs  1*254 
milligrams,  the  weight  in  milligrams  of  nitrogen  contained  in  that 
quantity  of  the  organic  compound  employed  for  the  analysis,  is 
obtained  by  multiplying  the  corrected  volume  (in  cubic  centimetres) 
by  1-254. 

Thus,  suppose  the  corrected  volume  of  nitrogen  =  50  c.c.,  then 
the  weight  of  nitrogen  would  be  50  x  1*254  =  6270  milligrams 
=  0*0627  gram. 

2.  Nitrogen  by  the  Soda-lime  Process. 

This  method  is  based  upon  the  fact  that  many  organic  substances 
containing  nitrogen  (see  p.  424),  when  strongly  heated  with  soda- 
lime,  give  up  their  nitrogen  in  combination  with  hydrogen  as 
ammonia.  By  estimating  the  ammonia  so  evolved,  the  weight  of 
nitrogen  can  be  determined. 

Epitome  of  Process. — The  organic  compound  is  mixed  with 
dry  granular  soda-lime  in  a  combustion-tube  closed  at  one  end,  and 
containing  a  short  layer  of  dry  oxalic  acid  at  its  closed  end  (in  order 
to  furnish  a  stream  of  hydrogen  wherewith  to  sweep  out  the  ammo- 
nia at  the  end  of  the  process).  The  evolved  ammonia  is  absorbed 
in  dilute  acid  contained  in  a  Will  and  Varrentrap's  bulb-tube,  which 
is  attached  to  the  combustion-tube  by  means  of  a  cork.  The 
ammonia  may  be  estimated  either  gravimetrically,  by  precipitation 
as  the  double  ammonium  platinic  chloride,  in  which  case  the  gas  is 
absorbed  in  dilute  hydrochloric  acid  (i  vol.  acid  to  4  vols. 
water),  and  the  estimation  carried  out  as  described  on  p.  251  ;  or 
it  may  be  determined  volumetrically  by  absorbing  it  in  an  excess 
of  standard  sulphuric  acid,  and  then  titrating  with  standard  sodium 
hydroxide.  The  latter  process  is  the  more  rapid. 

end  of  the  delivery  tube  beneath  a  separate  tube  (a  test-tube,  for  example) 
filled  with  potash  and  inverted  in  the  same  trough. 

At  the  conclusion  of  the  combustion,  the  measuring-tube  containing  the  gas 
is  withdrawn  from  the  mercury  trough  (by  slipping  a  small  dish  underneath  it) 
and  transferred  to  a  tall  cylinder  of  water.  The  mercury  and  the  remaining 
potash  flow  out,  and  are  replaced  by  water,  and  after  the  whole  has  assumed  a 
uniform  temperature,  the  tube  is  lowered  into  the  water  until  the  level  of  the 
liquid  inside  and  outside  are  the  same,  when  the  volume  is  read  off.  The 
temperature  of  the  water  and  the  barometric  pressure  are  noted.  In  correcting 
the  volume  to  N.T.P.,  the  full  value  of  the  pressure  due  to  aqueous  vapour 
must  be  taken  (compare  above). 


428  Ultimate  Organic  Analysis. 

Into  a  clean  dry  combustion-tube,  from  40  to  45  cms.  long,  is  in- 
troduced a  quantity  of  oxalic  acid  (previously  rendered  anhydrous  by 
being  heated  in  a  steam-oven),  sufficient  to  occupy  about  5  cms. 
Upon  this  is  added  about  the  same  quantity  of  dry  granular  soda- 
lime,  which  has  been  recently  heated  moderately  strongly  in  a 
porcelain  dish. 

A  quantity  of  soda-lime  (judged  to  be  sufficient  to  occupy  about 
10  cms.  of  the  tube)  is  powdered  in  a  dry  porcelain  mortar,  and  by 
means  of  a  small  clean  spatula  a  weighed  quantity  (from  0*3  to  0*5 
gram)  of  the  organic  compound  (urea  may  be  used  as  an  exercise) 
is  thoroughly  mixed  with  the  soda-lime.*  The  mixture  is  then 
transferred  to  a  sheet  of  clean  writing-paper,  and  carefully  poured 
into  the  tube  by  holding  the  paper  in  the  palm  of  one  hand  and 
bending  it  into  a  gutter.  The  mortar  is  then  rinsed  out  with  a 
little  more  powdered  soda-lime,  which  is  similarly  transferred  to 
the  tube.  The  tube  will  now  be  rather  more  than  half  full.  Granu- 
lar soda-lime  is  added  until  it  reaches  to  about  5  cms.  of  the  mouth, 
and  a  plug  of  asbestos  (previously  heated  to  redness)  is  inserted  to 
keep  the  materials  in  position.  The  tube  is  then  tapped  length- 
ways upon  the  table  (see  p.  425),  in  order  to  create  a  free  air-passage 
along  the  top  of  the  materials,  and  it  is  then  laid  in  the  furnace. 

A  measured  volume  of  normal  sulphuric  acid  (15  or  20  c.c., 
depending  upon  the  capacity  of  the  bulbs  (Fig.  94),  which  should 

contain  about  the 
volume  of  liquid  re- 
presented) is  intro- 
duced into  the  bulbs, 
which  are  then  con- 
nected to  the  tube  by 
a  tight  -  fitting  cork. 
The  portion  of  the 
tube  extending  back 
from  the  asbestos  plug, 
FlG-  94>  containing  soda  -  lime 

only,  is  first  heated  to  a  low  redness  ;  after  which  the  heating  is 
gradually  extended  along  until  the  whole  of  the  mixture  of  soda- 
lime  with  the  organic  compound  has  become  heated.  As  the 
heat  approaches  the  end  of  the  tube,  a  little  care  must  be  taken 
that  the  oxalic  acid  is  not  decomposed  before  its  time.  The  column 
of  pure  soda-lime  which  separates  the  mixture  should  serve  to 

*  Instead  of  mixing  the  materials  in  the  mortar,  the  process  may  be  carried 
out  within  the  tube  exactly  as  in  the  case  of  the  former  method  (p.  425). 


Nitrogen.  429 

protect  the  acid,  but  the  additional  precaution  may  be  taken  of 
inserting  in  the  furnace  a  small  screen,  made  of  asbestos  cardboard, 
before  beginning  the  operation. 

As  soon  as  the  evolution  of  ammonia  is  complete,  the  heating 
is  extended  to  the  oxalic  acid,  which  is  decomposed  in  the  presence 
of  excess  of  alkali  with  evolution  of  hydrogen — 

Na2C,O4  +  2NaHO  =  2Na2CO3  +  H2 

In  this  way,  the  ammonia  still  filling  the  tube  is  driven  out  into 
the  acid  in  the  bulbs.*  The  bulbs  are  then  disconnected  and  the 
contents  transferred  to  a  beaker,  the  bulbs  being  thoroughly  rinsed 
out  with  water.  The  excess  of  sulphuric  acid  present  is  titrated 
with  normal  sodium 'hydroxide  according  to  the  method  described 
on  p.  327. 

Each  cubic  centimetre  of  normal  acid  which  has  been  neutral- 
ised by  the  ammonia,  represents  o'ooiy  gram  of  ammonia,  or 
o'ooi4  gram  of  nitrogen. 

3.  Nitrogen  by  Kjeldalil's  Method. — This  process  depends 
upon  the  fact  that  many  nitrogenous  organic  compounds,  when 
heated  with  strong  sulphuric  acid,  have  their  nitrogen  converted 
into  ammonia,  which  at  once  unites  with  the  acid,  forming  am- 
monium sulphate.  From  this  salt  the  ammonia  is  afterwards 
expelled  by  means  of  caustic  soda,  and  estimated  by  one  of  the 
usual  methods. 

Epitome  of  Process.— The  nitrogenous  compound  is  digested 
with  strong  sulphuric  acid  in  a  small  round-bottom  flask  with  a 
long  neck,  potassium  sulphate  being  added  towards  the  end  of 
the  operation  in  order  to  raise  the  boiling-point  of  .the  liquid.  The 
mixture,  after  cooling,  is  diluted  with  water,  and  transferred  to  the 
apparatus  (Fig.  53,  p.  250),  where  it  is  treated  with  an  excess  of 
strong  sodium  hydroxide  solution,  and  the  ammonia  determined 
by  the  volumetric  method  given  on  p.  327. 

From  0*5  to  ro  gram  of  the  organic  compound  is  weighed  out 

*  In  this  process,  and  also  in  the  Dumas'  method  previously  described,  the 
combustion  may  be  conducted  in  a  tube  open  at  both  ends  (as  described  in 
the  estimation  of  carbon  and  hydrogen,  p.  415),  and  a  stream  of  carbon 
dioxide  or  of  hydrogen,  as  the  case  may  be,  passed  through  from  a  reservoir 
or  generating  apparatus  outside.  This  modification  of  the  processes,  how- 
ever, is  not  found  to  yield  better  results ;  and  as  in  both  cases  the  compound 
must  be  mixed  with  the  contents  of  the  tube,  and  cannot  be  introduced  in  a 
boat,  the  tube  has  to  be  dismounted  and  recharged  for  each  combustion.  In 
the  carbon  and  hydrogen  combustion,  the  open-tube  method  enables  a  series 
of  determinations  to  be  made  in  rapid  succession  without  dismantling  the 
apparatus,  but  in  a  nitrogen  determination  this  is  not  the  case,  and  it  not 
only  involves  the  use  of  unnecessary  apparatus,  but  introduces  manipulative 
difficulties  in  filling  the  combustion-tube. 

2  F 


43°  Ultimate  Organic  Analysis. 

into  the  digestion-flask,  and  20  c.c.  of  strong  sulphuric  acid  added 
to  it.*  The  flask  is  then  supported  in  an  inclined  position,  and 
heated  by  means  of  a  small  flame  nearly  to  the  boiling-point  until 
the  mixture  ceases  to  froth.  The  temperature  is  then  raised,  and 
the  mixture  (usually  dark-coloured  or  black)  boiled  for  about  15 
minutes,  carefully  guarding  against  frothing.  Ten  to  12  grams 
of  potassium  sulphate  are  then  added,  and  the  mixture  allowed  to 
continue  gently  boiling  until  it  is  nearly  colourless  ;  after  which 
the  flame  is  removed,  and  the  liquid  allowed  to  cool.  A  little  water 
is  added,  and  the  contents  of  the  flask  thoroughly  rinsed  out  into 
the  distilling  apparatus  shown  on  p.  250.  It  is  there  treated  with 
caustic  soda,  and  the  ammonia  distilled  into  an  excess  of  standard 
sulphuric  acid  as  described  on  p.  327. 

III.  Estimation  of  Chlorine  t  in  Organic  Compounds. 

Epitome  of  Process  (Carius*  method).—  A  weighed  quantity 
of  the  substance,  contained  in  a  thin  tube  or  bulb,  is  introduced 
into  a  piece  of  stout  combustion-tube,  sealed  at  one  end,  along 
with  a  few  crystals  of  silver  nitrate  and  a  quantity  of  strong  nitric 
acid.  The  open  end  of  the  tube  is  then  drawn  out  and  sealed. 
The  thin  tube  containing  the  compound  is  then  broken,  and  the 
combustion-tube  is  placed  in  a  hot-air  bath,  specially  constructed 
for  the  purpose,  and  heated  to  a  temperature  necessary  to  ensure 
complete  decomposition  of  the  compound,  for  two  or  three  hours. 
In  this  way  the  carbon  and  hydrogen  are  converted  into  carbon 
dioxide  and  water,  while  the  liberated  halogen  combines  with  the 
silver,  present  as  nitrate,  and  forms  silver  chloride.  The  tube  is 
afterwards  opened  and  the  contents  washed  out,  and  the  silver 
chloride  filtered,  dried,  and  weighed  in  the  usual  way. 

A  piece  of  combustion-tube  about  60  cms.  long,  not  too  wide  in 
the  bore,  and  with  tolerably  thick  walls,  is  sealed  at  one  end  before 
the  blowpipe,  care  being  taken  not  to  allow  the  end  to  be  thinned 
in  the  process.  The  organic  compound  is  introduced  into  a  thin 
glass  tube,  the  end  of  which  has  been  slightly  enlarged  before 
the  blowpipe  into  a  still  thinner  bulb  (a,  Fig.  95).  This  little 
bulb-tube  is  first  weighed  ;  about  0*3  to  0-5  gram  of  (he  substance  is 
introduced,  and  the  tube  drawn  off  and  sealed  by  means  of  a  fine 
pointed  blowpipe-flame  at  a1.  The  two  portions  a  and  b  are  then 
together  weighed.  The  difference  is  the  weight  of  the  substance. 

*  As  the  sulphuric  acid  of  commerce  is  liable  to  contain  ammonium 
sulphate  as  an  impurity,  a  blank  experiment  should  be  made  with  the  sample 
of  acid  which  is  to  be  used  ;  and  the  weight  of  nitrogen  per  cubic  centimetre 
of  the  acid  which  is  represented  by  this  impurity,  must  be  deducted  from  the 
result  obtained  in  subsequent  analyses. 

•f  The  process  is  equally  applicable  in  the  cases  of  bromine  and  iodine. 


Chlorine. 


431 


FIG.  95. 


The  portion  b  is  left  upon  the  scale-pan,  while  the  bulb  is  carefully 
introduced  into  the  combustion-tube,  into  which  a  few  crystals  (i 
or  2  grams)  of  silver  nitrate  have  previously  been  dropped.  Fuming 
nitric  acid  (free  from 
chlorine)  is  then  added 
in  quantity  sufficient 
to  occupy  about  5  cms. 
of  the  tube.  The  tube 
is  then  held  in  an  in- 
clined position  before 
the  blowpipe,  and 
drawn  out  into  a  thick 
capillary.  This  is  done 
by  heating  the  tube  in 
the  flame,  revolving  it 
slowly,  until  the  soft- 
ened walls  thicken  and 
almost  fall  together, 
assuming  the  condition  shown  in  C,  Fig.  95.  The  glass  is  then 
gently  pulled  out,  when  the  thick  capillary  shown  at  D  will  be 
obtained  ;  the  end  is  then  sealed  in  the  flame.  When  the  tube  has 
become  quite  cold,  the  little  bulb  within  is  broken  by  shaking  the 
tube  in  such  a  way  as  to  cause  the  bulb  to  strike  against  the  walls, 
but  without  allowing  any  of  the  contents  to  get  into  the  capillary 
end.  The  tube  is  then  laid  in  the  air-bath  and  gradually  heated 
up  to  a  temperature  of  1 50°  to  200°. 

The  air-bath  consists  of  an  oblong  iron  chamber  or  box,  with 
a  number  of  iron  tubes  closed  at  one  end  passing  through  length- 
ways ;  the  tubes  are  screwed  into  the  ends  of  the  box,  and  fixed 
in  a  slightly  inclined  position,  so  that  when  the  glass  tubes  are 
pushed  in,  the  liquid  contents  will  be  kept  at  one  end.  The  whole 
is  supported  on  an  iron  stand,  and  heated  by  a  row  of  small  flames 
placed  beneath.  A  thermometer,  passing  through  a  small  hole  in 
the  top,  registers  the  temperature.  As  great  pressure  is  often 
developed  in  the  glass  tubes,  it  is  desirable  to  place  an  iron  screen 
or  a  brick  opposite  the  end  of  the  bath,  so  that,  in  the  event  of  a 
tube  bursting,  the  glass  shall  be  prevented  from  flying  about. 
After  about  two  hours'  heating,  the  gas  is  extinguished,  and  the 
glass  tube  allowed  to  become  quite  cold  before  being  removed. 

It  must  never  be  forgotten  that  the  tube  is  now  a  somewhat 
dangerous  object,  and  it  must  therefore  be  handled  with  caution. 
Such  a  tube  should  never  be  left  lying  about,  as  they  sometimes 


432  Ultimate  Organic  Analysis. 

burst  with  great  violence,  without  any  apparent  reason,  even  many 
hours  after  they  have  become  cold. 

It  is  carefully  withdrawn  from  the  bath,  and  immediately  en- 
veloped in  a  cloth,  with  the  end  of  the  capillary  tube  only  projecting, 
and  keeping  it  all  the  time  in  such  a  position  that  the  liquid  does 
not  run  up  to  the  narrow  end.  Should  any  liquid  have  got  into 
the  end,  it  is  carefully  driven  out  by  gently  brushing  a  flame  along 
that  part  of  the  tube. 

The  tube  is  then  opened  by  cautiously  introducing  the  tip  of 
the  capillary  into  a  blowpipe-flame,  so  as  to  gradually  soften  the 
glass,  which  is  then  blown  out  by  the  outrushing  gases  from  the  tube. 

On  no  account  whatever  must  the  tube  be  opened  by  scratching 
the  capillary  with  a  file  and  breaking  off  the  point,  as  this  is 
extremely  likely  to  cause  the  tube  to  burst. 

When  the  gas  has  escaped,  the  tube  is  cut  at  a  point  just  below 
the  narrowed  portion  by  first  making  a  scratch  with  a  sharp  file, 
and  then  touching  the  end  of  the  scratch  with  the  red-hot  end  of 
a  piece  of  glass  or  wire.  The  contents  of  the  tube  are  then  trans- 
ferred to  a  beaker,  the  tube  and  the  cut-off  piece  being  thoroughly 
rinsed  into  the  beaker. 

Besides  the  silver  chloride,  there  will  be  the  remains  of  the 
glass  bulb-tube.  As  the  latter  will  probably  have  been  broken  at 
the  blown-out  and  thinner  end,  it  is  generally  possible  to  pick  out, 
with  a  platinum  wire,  the  greater  part  of  the  glass  in  a  single  large 
piece.  This  is  rinsed  thoroughly,  and  carefully  placed  on  a  piece 
of  filter-paper  in  a  steam-oven  to  dry.  It  is  then  conveyed  to  the 
balance,  and  weighed  along  with  the  drawn-out  end  which  has 
been  left  there.  The  difference  between  this  weight  and  the  original 
weight  of  the  empty  bulb-tube,  will  represent  the  weight  of  the 
fragments  of  glass  still  remaining  admixed  with  the  silver  chloride, 
and  which  therefore  must  afterwards  be  deducted  from  the  weight 
finally  obtained. 

To  the  acid  liquid  sufficient  sodium  hydroxide  (free  from 
chloride)  is  added  to  neutralise  most  of  the  acid,  and  the  mixture 
is  then  filtered,  every  particle  of  the  broken  glass,  as  well  as  the 
silver  chloride,  being  carefully  brought  on  to  the  filter.  The  pre- 
cipitate is  washed,  dried,  and  treated  in  the  usual  way. 

From  the  weight  of  silver  chloride  (after  deducting  the  weight 
of  the  fragments  of  glass)  the  percentage  of  chlorine  in  the  com- 
pound is  calculated. 

IV.  Estimation  of  Sulphur  and  Phosphorus  in  Organic 
Compounds. — The  oxidation  of  the  organic  compound  by  means 


Sulphur  and  Phosphorus.  433 

of  strong  nitric  acid  in  a  sealed  tube  converts  sulphur  into  sulphuric 
acid,  and  phosphorus  into  phosphoric  acid.  In  either  case,  there- 
fore, the  process  is  carried  out  as  in  the  estimation  of  chlorine, 
but  without  the  silver  nitrate. 

In  the  case  of  sulphur,  the  contents  of  the  tube  are  transferred 
to  a  dish,  and  the  nitric  acid  expelled  by  evaporating  the  mixture 
nearly  down  to  dry  ness.  Dilute  hydrochloric  acid  is  added,  and 
the  sulphuric  acid  precipitated  as  barium  sulphate  in  the  usual 
way  (p.  258).  When  phosphorus  is  being  estimated,  the  acid 
liquid  is  neutralised  with  ammonia,  and  the  phosphoric  acid  pre- 
cipitated as  ammonium  magnesium  phosphate  (p.  264). 


SECTION    II. 


MISCELLANEOUS  PHYSICO-CHEMICAL  DETERMINATIONS. 

SPECIFIC  gravity  of  solids  and  liquids. 
Boiling-point. 
Melting-point. 
Vapour  density. 

(i)  Determination  of  the  Specific  Gravity  of  a  Solid.* 

(a)  Solid  Substances  which  are  not  acted  upon  by  Water. 
Principle. — When  a  body  is  weighed  suspended  in  a  liquid,  its 
weight  is  diminished  by  the  weight  of 
the  liquid  displaced.     Hence,  therefore, 
if  the  body  is  weighed  first  in  air,  and 
^\^  again  when  immersed  in  water,  the  dif- 

^   ^5v  ference  between  the  two  weights  is  the 

>5^  weight  of  the  water  displaced  by  the 
body  ;  that  is,  the  volume  of  water  which 
is  equal  to  the  volume  of  the  body. 

EXAMPLE. — A  fragment  of  quartz  or 
flint  or  other  mineral,  about  the  size  of  a 
walnut,  is  weighed  in  the  ordinary  way. 
It  is  then  suspended  by  means  of  a 
thread  of  silk,  or  a  horsehair,  from  the 
beam  of  the  balance,  in  such  a  manner 

*  The  specific  gravity  of  a  solid  or  liquid 
substance  is  the  ratio  between  the  weight  of  a 
given  volume  of  the  substance  and  the  weight 
of  the  same  volume  of  water  at  its  point  of 
maximum  density  (taken  as  4°  C. ).  It  is, 
therefore,  obtained  by  dividing  the  former 
weight  by  the  latter.  As,  however,  it  is  more 
difficult  to  maintain  a  uniform  temperature  at 
4  than  at  the  common  atmospheric  tempera- 
ture, the  operations  are  usually  conducted  at 
FIG.  96-  15-5- 


Specific  Gravity.  435 

that  it  will  reach  to  about  the  middle  of  a  small  beaker  supported 
upon  a  wooden  bridge  which  stands  over  the  pan,  as  in  Fig.  96, 
leaving  the  latter  free  to  move.  The  beaker  is  nearly  filled  with 
water  at  /5°'5,  and  any  air-bubbles  adhering  to  the  solid  substance 
are  carefully  removed  by  means  of  a  feather  or  small  camel's-hair 
brush.*  In  this  position  the  substance  is  again  weighed. t 

If  a  =  the  weight  in  air,  and  b  —  weight  in  water — 
then  a  —  b  =  the  weight  of  water  displaced 

therefore — — ,  =  specific  gravity  of  the  substance. 

(b)  Solid  Substances  which  are  acted  upon  by  Water. — When 
the  solid  is  dissolved,  or  otherwise  acted  upon  by  water,  some  other 
liquid  of  known  specific  gravity  is  employed  which  exerts  no  action 
upon  the  body. 

If  x  =  density  of  the  solid  as  compared  with  the  liquid, 
and_y  =  specific  gravity  of  the  liquid — 

(sp.  gr.  water) 

then       i       :  y  '. '.  x  \  specific  gravity  of  solid  compared  with  water. 

EXAMPLE. — A  crystal  of  copper  sulphate  (or  alum,  nitre,  sodium 
thiosulphate,  etc.)  is  first  weighed  in  the  usual  manner.  It  is  then 
suspended  in  a  beaker  containing  benzene  (C6H6,  sp.  gr.  0-885),  and 
again  weighed,  care  being  taken,  as  in  the  former  example,  to 
remove  all  air-bubbles. 

If  a  —  weight  in  air,  and  b  =  weight  in  benzene — 

then  —  —    =  density  as  compared  with  benzene  =  x 
and  x  x  y  =  specific  gravity  compared  with  water. 

(<:)  Powdered  Substances. — When  the  substance  is  in  a  powdered 
condition,  the  weight  of  water  which  it  displaces  is  ascertained  by 
weighing  it  in  a  small  bottle  which  is  filled  up  with  water.  The 
weight  of  the  empty  bottle  is  known,  and  the  weight  of  water  it 
contains  when  full  is  also  known. 

*  A  porous  substance,  like  charcoal,  for  instance,  must  be  first  placed  in 
water  beneath  an  air-pump  receiver  and  exhausted. 

f  If  the  substance  is  lighter  \hnn.  water,  a  weight  is  attached  to  it  in  order  to 
sink  it.     The  weight  of  this  singer  \\hen  suspended  in  water  is  first  ascertained. 
Then,  if  a  =  weight  of  substance  in  air, 

b  =  weight  of  substance  and  sinker  together  in  water, 
c  —  weight  of  sinker  in  water — 

then    7v =  specific  gravity  of  the  substance 


436  Pliysico-chemical  Determinations. 

If  a  —  weight  of  powder  in.  air, 
b  —  weight  of  powder  and  water, 
c  —  weight  of  water — 

then  -      ,.          =  specific  gravity  of  the  powder 

A  small  bottle  with  a  perforated  glass  stopper  (known  as  a 
specific-gravity  bottle)  is  exactly  weighed  in  a  perfectly  clean  and 
dry  condition.  It  is  then  filled  with  water  at  I50<5,  and  the  stopper 
inserted.  The  small  quantity  of  water  which  is  thereby  extruded 
through  the  bored  stopper  is  wiped  off  with  a  clean  cloth,  and  the 
bottle  quickly  weighed.  (As  sold,  these  bottles  are  marked  with 
the  weight  of  water  they  are  supposed  to  hold  at  this  temperature.) 
From  this  the  weight  of  water  it  contains  (c  in  the  above  formula) 
is  ascertained  once  for  all. 

Into  the  dry  bottle  a  quantity  of  the  powder  is  placed,  and  the 
bottle  weighed.  From  this  the  weight  of  powder  in  air  (a)  is  found 
by  deducting  the  known  weight  of  the  empty  bottle.  The  bottle 
containing  the  powder  is  then  filled  with  water  at  i5°'5,  being 
gently  tapped  to  disentangle  any  air-bubbles.  The  stopper  is  in- 
serted, and  after  carefully  wiping  off  the  overflow  the  bottle  is  again 
weighed.  By  subtracting  the  weight  of  the  empty  bottle,  this  gives 
the  weight  of  powder  +  water  (6).  From  these  data  the  specific 
gravity  is  calculated  by  the  above  formula. 

(2)  Determination  of  the  Specific  Gravity  of  a  Liquid. 
— This  may  be  ascertained  by  means  of  the  specific-gravity  bottle 
described   above.      It   is  filled  with   the  liquid   in  question  at  a 
temperature  of  15°'$  and  weighed.     Deducting  the  weight  of  the 
empty  bottle  gives  the  weight  of  the  liquid.     The  weight  of  the 
same  volume  of  water  at  the  same  temperature  being  already  known, 
the  specific  gravity  of  the  liquid  is  at  once  ascertained  by  dividing 
the  former  by  the  latter. 

Instead  of  the  specific-gravity  bottle,  a  thin  glass  U-tube,  drawn 
out  at  each  end  to  a  capillary,  may  be  employed.  The  liquid  is  drawn 
into  the  tube  by  dipping  one  extremity  into  the  bottle  containing  it, 
and  applying  a  gentle  suction  to  a  piece  of  rubber  tube  attached  to 
the  other  end.  The  liquid  is  then  brought  to  the  temperature  of 
15°  5  by  immersing  it  in  a  beaker  of  water  at  this  temperature, 
after  which  it  is  carefully  wiped  and  weighed. 

(3)  Determination  of  Boiling-point.— The  liquid  is  gently 
heated  in  a  tubulated  distillmg-flask  (a  "Wurtz"  flask)  fitted  with 
a  cork,  through  which  is  inserted  a  thermometer,  the  bulb  of  the 


Boiling-point.  437 

latter  being  a  little  above  the  surface  of  the  liquid.  The  flask 
may  be  connected  to  a  condenser  in  order  to  recover  the  portion  of 
the  liquid  which  is  vaporised  during  the  process.  A  fragment  of 
scrumpled  platinum  foil  placed  in  the  flask  causes  the  liquid  to 
boil  more  regularly. 

The  point  at  which  the  mercury  in  the  thermometer  remains 
constant  is  noted,  and  where  only  moderate  exactness  is  required 
this  temperature  may  be  taken  as  the  boiling-point. 

In  more  exact  determinations,  either  a  correction  must  be  made 
for  the  cooling  of  the  mercury  in  such  portions  of  the  column  as 
extend  above  the  cork,  or  else  the  operation  must  be  performed  in 
such  an  apparatus  as  to  ensure  the  equal  heating  of  the  entire  thread 
of  mercury. 

In  the  former  case  a  second  thermometer  is  lashed  to  the  first 
outside  the  flask,  in  such  a  manner  that  the  bulb  lies  midway 
between  the  cork  and  the  highest  point  to  which  the  mercury 
reaches,  so  as  to  record  the  average  temperature  of  that  portion  of 
the  mercury  thread  which  extends  above  the  cork.  The  correction 
which  is  added  to  the  observed  temperature  is  then  made  by  means 
of  the  following  formula  : — 

n(t  —  I')  x  0-000143 

where n  —  number  of  degrees  over  which  the  mercury  is  not  exposed 

to  the  heated  vapour  ; 
/  —  observed  temperature  ; 
/  =  mean  temperature  of  the   outside  mercury  (given   by 

the  second  thermometer)  ;  and 
0000143  =  apparent  coefficient  of  expansion  of  mercury  in  glass. 

More  exact  determinations  of  boiling-points  are  made  by  means 
of  the  apparatus  shown  in  Fig.  97.  The  vapour  from  the  boiling 
liquid  passes  up  the  central  tube,  and  then  down  through  the 
annular  space  between  the  inner  and  outer  tubes  before  it  finally 
issues  through  the  lateral  branch  tube.  This  side  tube,  which 
extends  for  a  considerable  length,  is  connected  to  a  condenser. 

In  this  apparatus,  not  only  is  the  whole  of  the  mercury  thread 
of  the  thermometer  exposed  to  the  hot  vapour,  but  the  tube  through 
which  it  passes  is  itself  surrounded  by  a  jacket  of  the  same  vapour. 
In  this  way  cooling  due  to  radiation  is  reduced  to  a  minimum.* 

*  It  will  be  obvious  that  when  a  boiling-point  is  to  be  determined  with 
sufficient  exactness  to  necessitate  the  use  of  this  apparatus,  a  thermometer 
which  has  been  standardised  must  be  employed,  and  the  instrument  must  be 
one  having  a  sufficiently  open  scale  to  render  it  possible  to  read  small  fractions 
of  degrees. 


438 


Physico-chem  ical  Determ {nations. 


It  must  be  remembered  that  the  temperature  at  which  a 
liquid  boils  is  influenced  by  pressure  ;  thus  when  we  say  that  the 
boiling-point  of  water  is  100°,  it  is  understood  that  the  liquid  boils 
at  this  temperature  under  normal  atmospheric 
pressure.  When  making  a  boiling-point 
determination,  therefore,  the  barometric  pres- 
sure at  the  time  must  be  noted.  If  this  is 
greater  or  less  than  760  mm.,  then  either  the 
observed  boiling-point  is  recorded  as  having 
been  taken  under  that  particular  pressure,  or 
a  correction  for  pressure  must  be  introduced. 
There  is,  however,  no  formula  of  general 
application  by  means  of  which  this  correction 
can  be  made,  for  the  reason  that  the  extent 
to  which  the  boiling-point  is  influenced  by 
variations  of  pressure  is  different  with  different 
liquids.  If,  therefore,  the  liquid  under  exami- 
nation is  an  unknown  substance,  its  exact 
normal  boiling-point  (i.e.  its  boiling-point  at 
normal  pressure)  can  only  be  ascertained  by 
making  the  determination  at  that  actual 
pressure. 

The  extent  to  which  liquids  differ  from 
each  other  with  respect  to  the  effect  of  pres- 
sure upon  their  boiling-points  is,  however, 
very  slight,  so  that  an  approximate  correction 
may  be  introduced  by  regarding  them  as  all 
behaving  as  water  does. 
For  each  millimetre  pressure  (in  the  neighbourhood  of  760 
mm.)  the  boiling-point  of  water  varies  0-0375°  5  or)  m  other  words, 
the  boiling-point  is  lowered  o-i°  for  a  decrease  of  pressure  amount- 
ing to  2-68  mm.,  and  raised  by  the  same  faction  by  a  similar 
increase  of  pressure. 

(4)  Determination  of  Melting-point. — Provided  a  sub- 
stance melts  without  undergoing  decomposition,  the  temperature 
at  which  it  fuses  when  in  a  state  of  purity  is  constant,  and  is 
termed  its  melting-point.  This  is  determined  by  introducing  the 
solid  substance  into  a  thin  glass  tube  drawn  down  to  a  capillary 
end  which  is  sealed.  This  little  tube  is  lashed  to  the  bulb  of  a 
thermometer  either  by  small  rubber  rings  (cut  from  ordinary  rubber 
tube)  or  by  wire,  depending  upon  the  particular  liquid  to  be 
employed  in  the  bath  in  which  it  is  to  be  heated.  Thus  if  a  bath 


FIG.  97. 


Vapour-density  Determination. 


439 


of  strong  sulphuric  acid  is  used,  platinum  wire  should  be  employed. 
Fig.  98  shows  the  tube  attached  to  the  thermometer.  The  ther- 
mometer is  then  supported  in  a  clamp,  and  lowered  into  a  tall 
narrow  beaker  containing  a  liquid  which  may  be  heated 
well  above  the  melting-point  of  the  substance,  without 
boiling.  A  small  flame  is  placed  beneath  the  beaker, 
and  the  liquid  is  kept  constantly  mixed  by  means  of 
a  stirrer,  consisting  of  a  glass  rod  bent  in  the  form  of 
a  ring  at  one  end.  This  is  moved  up  and  down  in 
the  beaker,  and  the  moment  the  solid  substance  melts, 
the  temperature  is  read  off. 

The  thermometer  is  then  raised  out  of  the  bath, 
which  is  allowed  to  cool  somewhat,  and  the  operation 
repeated.  With  some  solids,  such  as  fats,  etc.,  the 
repetition  of  the  determination  must  be  made  with  a 
fresh  specimen,  as  each  time  they  are  melted  they 
undergo  a  slight  decomposition  or  change  which  lowers 
the  melting-point.  If  the  substance  is  translucent, 
or  is  one  in  which  from  other  causes  it  is  difficult 
to  observe  the  exact  fusion  point,  the  capillary  tube 
may  be  open  at  the  end,  and  the  substance  introduced 
by  dipping  the  tube  into  a  small  quantity  of  the  melted 
compound.  The  liquid  rises  by  capillarity  into  the 
tube,  which  is  then  removed  and  the  substance  allowed 
to  re-solidify.  When  the  tube  so  charged  is  intro- 
duced as  before  into  the  bath,  the  moment  the  sub- 
stance melts  it  is  driven  up  the  tube  by  the  entrance  of 
the  liquid  from  the  bath,  and  the  temperature  at  which 
this  movement  is  observed  is  noted.  As  in  the  boiling- 
point  determination,  the  same  correction  must  be  made 
for  the  portion  of  the  mercury  thread  which  projects 
out  of  the  bath. 

(5)  Determination  of  Vapour  Density.— The  density  of 
a  vapour  is  the  ratio  between  the  weight  of  a  given  volume  of  that 
vapour,  and  the  weight  of  the  same  volume  of  another  gas  (under 
the  same  conditions  of  temperature  and  pressure)  which  is  taken 
as  the  standard  or  unit.  The  standard  usually  adopted  is  hydrogen; 
therefore  the  vapour  density  of  a  substance  is  in  reality  the  specific 
gravity  of  the  vapour  of  that  substance  compared  with  hydrogen. 

Vapour  Density  by  Victor  Meyer's  Method. — Of  the 
various  methods  which  have  been  devised  for  determining  vapour 
densities,  that  of  Victor  Meyer  is  the  one  most  suitable  for  general 


FIG.  98. 


440 


Physico-chemical  Determinations. 


laboratory  practice,  where  a  very  high  degree  of  accuracy  is  not 
desired.*   The  apparatus  consists  of  a  bulb-tube,  b  (Fig.  99),  with  a 

long,  rather  narrow  stem,  upon 
'which  near  the  top  there  is  a  branch 
tube  of  narrow  bore.  This  bulb- 
tube  is  placed  in  a  wider  tube,  in 
which  a  suitable  liquid  may  be 
boiled  in  order  that  the  bulb-tube 
may  be  heated  by  the  vapour.  A 
weighed  quantity  of  the  liquid 
whose  vapour  density  is  to  be  taken 
is  introduced  into  the  hot  bulb- 
tube,  which  is  then  immediately 
closed  with  a  cork.  The  liquid  in 
vaporising  displaces  a  volume  of 
air  equal  to  the  volume  of  the 
vapour,  and  the  air  so  expelled  is 
collected  over  water  in  the  gra- 
duated tube,  under  the  ordinary 
atmospheric  conditions  of  tempera- 
ture and  pressure,  which  are  noted. 
The  volume  in  cubic  centimetres 
so  measured  is  reduced  to  N.T.P., 
and  the  corrected  number  of  cubic 
centimetres  when  multiplied  by 
0-0000896  (the  weight  in  grams  of 
i  c.c.  of  hydrogen)  gives  the  weight 
of  an  equal  volume  of  hydrogen. 

For  practice  the  vapour  den- 
sity of  carbon  disulphide  may  be 
taken. f  A  little  plug  of  glass  wool  or  asbestos  is  pushed  down  to 
the  bottom  of  the  bulb-tube  ^,  which  is  then  placed  in  the  outer 
tube,  into  which  a  small  quantity  of  water  has  been  introduced. 
The  tube  b  must  be  perfectly  dry  inside.  The  top  is  closed  with 
a  good  cork,  and  the  water  boiled.  As  the  temperature  rises  and 
the  air  in  b  expands,  the  excess  bubbles  out  through  the  water  in 
the  small  trough.  When  the  temperature  is  constant,  and  no  more 

*  For  descriptions  of  other  methods,  as  Dumas',  Hofmann's,  etc.,  the 
student  is  referred  to  text-books  on  Physics. 

t  The  specimen  should  be  as  pure  as  possible.  It  may  be  distilled,  and 
that  portion  of  the  distillate  which  passes  over  at  46°  collected  separately  for 
the  determination. 


FIG.  99. 


Vapour-density  Determination. 


441 


air  is  expelled,  the  graduated  measuring-tube  filled  with  water  is 
brought  over  the  end  of  the  bent  delivery  tube.  The  apparatus  is 
now  ready  for  the  introduction  of  the  liquid.*  This  is  contained 
in  a  tiny  stoppered  bottle,  or  in  a  fine  glass  tube  sealed  at  one  end 
and  drawn  to  a  capillary  at  the  other  (shown  the  actual  size  at  a 
and  b,  Fig.  100).  The  bottle  or  tube  is  weighed  first  empty,  and 
again  when  filled  with  the  liquid  (the  tube  is 
filled  by  gently  heating  in  a  flame,  and  then 
dipping  the  end  into  the  liquid  ;  and  the 
capillary  end  is  left  unsealed).  About  OT 
gram  is  a  suitable  weight  of  liquid  to  employ. 

The  cork  is  removed  from  the  bulb-tube, 
and  the  little  bottle  (or  tube),  held  in  a  pair 
of  forceps,  is  dropped  in,  and  the  cork  quickly 
replaced.  If  the  bottle  is  used,  the  stopper  is 
just  loosened  before  dropping  it  in  ;  if  the 
tube  is  employed,  it  is  dropped  in  capillary 
end  downwards.  The  pad  of  glass  wool  or 
asbestos  prevents  the  bulb  from  being  frac- 
tured by  the  falling  bottle.  The  liquid  imme- 
diately vaporises  and  expels  air  from  the 
apparatus,  which  collects  in  the  graduated  tube.  When  no  more 
bubbles  escape,  the  cork  is  removed. 

The  measuring-tube  is  then  transferred  to  a  large  beaker  filled 
with  water  at  the  temperature  of  the  room,  the  transference  being 
done  by  slipping  a  small  porcelain  dish  under  the  tube  as  it  stands 
in  the  water-trough.  The  tube  is  depressed  in  the  water  until  the 

*  Other  liquids  than  water  are  used  in  the  outer  tube  or  bath,  when  a 
higher  temperature  is  required  to  vaporise  the  substance  to  be  examined.  It 
is  only  necessary  that  the  temperature  of  this  bath  should  be  considerably 
higher  than  the  boiling-point  of  the  liquid  to  be  vaporised,  and  that  it  should 
remain  constant  throughout  the  operation.  It  is  not  necessary  to  know  what 
the  actual  temperature  is.  This  fact  often  presents  a  difficulty  to  the  young 
student,  which  is  expressed  by  the  constantly  repeated  question,  "If  the 
volume  of  gas  or  vapour  is  increased  by  rise  of  temperature,  then  surely  the 
hotter  the  tube  is,  the  greater  the  volume  of  vapour  which  will  be  produced 
when  the  liquid  is  introduced,  and  therefore  a  correspondingly  greater  volume 
of  air  will  be  displaced  from  the  apparatus ;  how  is  it,  therefore,  that  it  is  not 
necessary  to  take  the  temperature  of  the  bath  ?"  This  difficulty  vanishes  when 
the  fact  is  taken  into  account  that,  whatever  is  the  temperature  of  the  bath 
when  the  volatile  liquid  is  introduced,  the  air  within  has  already  been  heated  to 
that  temperature,  and  proportionately  expanded  or  rarefied.  Consequently 
at  the  higher  temperature,  although  the  volume  occupied  by  the  vapour  is 
greater,  and  therefore  the  volume  of  air  expelled  is  greater,  it  is  only  more 
rarefied  air,  which  when  cooled  by  collection  over  the  water,  is  reduced  to  the 
same  volume  as  the  lesser  bulk  of  less  expanded  air  which  would  be  displaced 
by  the  less  expanded  vapour  at  the  lower  temperature. 


44  2  Physico-chemical  Determinations. 

level  of  the  liquid  inside  and  outside  is  the  same,  when  the  volume 
is  read  off.  The  temperature  of  the  water  is  taken,  and  height  of 
the  barometer  is  read. 

The  gas  being  saturated  with  aqueous  vapour,  the  volume  is 
reduced  to  N.T  P.  by  the  formula  given  on  p.  388. 

The  corrected  volume  so  obtained  is  the  number  of  cubic 
centimetres  which  the  vapour  of  the  carbon  disulphide  employed 
would  occupy  if  it  could  exist  under  N.T. P. 

Multiplying  this  by  0-0000896  gives  the  weight  in  grams  of  an 
equal  volume  of  hydrogen  at  N.T.P.  ;  and  by  dividing  the  weight 
of  carbon  disulphide  employed  by  this  weight  of  hydrogen,  the 
vapour  density  of  carbon  disulphide  is  obtained. 

6.  Determination  of  Molecular  Weight  by  the  "Freez- 
ing-point Method." — When  pure  water  is  cooled  under  ordinary 
conditions,  it  freezes  at  o°  C.  ;  but  if  the  water  be  rendered  impure 
by  the  presence  of  dissolved  salts  or  other  substances,  it  becomes 
necessary  to  cool  it  below  o°  before  ice  begins  to  form.  Thus, 
water  containing  I  per  cent,  of  sodium  chloride  in  solution  will 
require  to  be  cooled  to  —  o'6°  before  it  begins  to  freeze.  It  is  a 
familiar  fact  that  a  lower  degree  of  cold  is  necessary  to  freeze  sea- 
water  than  fresh. 

When  such  aqueous  solutions  are  cooled,  the  solid  which  begins 
to  separate  out  is  pure  ice,  and  not  the  substance  which  is  in  solu- 
tion in  the  water,  provided  the  solution  is  sufficiently  dilute. 

What  is  true  of  water  in  this  respect  is  true  also  of  other  liquids 
which  are  capable  of  being  solidified.  Thus,  the  freezing-point  of 
pure  benzene  is  6°,  but  if  a  small  quantity  of  benzoic  acid  (a  solid) 
or  aniline  (a  liquid)  be  dissolved  in  the  benzene,  then  it  will  be 
found  that  this  solution  will  require  to  be  cooled  below  6°  before 
the  benzene  begins  to  freeze  ;  and,  just  as  in  the  case  of  water,  it  is 
the  solvent  alone  which  freezes  out  of  the  solution. 

The  discovery  that  a  quantitative  relation  existed  between  the 
lowering  of  the  freezing-point  of  a  liquid  and  the  amount  of  dissolved 
substance  present  was  made  by  Blagden  as  long  ago  as  1788,  who 
formulated  the  law  that  the  depression  of  the  freezing-point  of 
aqueous  solutions  of  the  same  substance  was  proportional  to  the 
amount  of  substance  dissolved. 

It  was  not,  however,  until  the  year  1871  that  the  important 
observation  was  made  by  Coppet  that,  in  the  case  of  certain  chemi- 
cally allied  substances,  if  quantities  which  wet  proportional  to  the 
molecular  weights  were  dissolved  in  equal  volumes  of  the  same 
solvent,  the  freezing-point  was  depressed  to  the  same  extent ;  or, 


Molecular  Weight  by  Freezing-point  Method. 


in  other  words,  the  solutions  would  have  the  same  solidifying- 
point. 

Upon  these  fundamental  principles  is  based  the  process  developed 
by  Raoult,  and  known  as  Raoult's  method  (or  the  cryoscopic  method) 
for  determining  molecular  weights. 

The  extent  to  which  the  freezing-point  of  a  liquid  would  be 
depressed  by  dissolving  in  100  grams  of  it  a  gramme-molecule  (i.e. 
a  weight  in  grams  equal  to  the  molecular  weight)  of  any  substance, 
is  spoken  of  as  the  molecular  lowering,  or  molecular  depression^  of 
the  freezing-point  of  that  liquid  ;  and  with  certain  qualifications, 
this  is  found,  for  one  and  the  same  solvent,  to  approximate  to  a 
constant*  Thus  when  water  is  the  solvent,  the  molecular  lowering 
of  the  freezing-point  is  about  i8'5.  For  acetic  acid  the  constant  is 
about  39,  while  for  benzene  the  value  is  about  49. 

For  example,  a  gramme-molecule  of  acetamide  dissolved  in  100 
grams  of  water  would  depress  the  freezing-point  17  '8°,  while  the 
same  weight  dissolved  in  100  grams  of  acetic  acid  lowers  the 
freezing-point  of  this  solvent  36'  i°. 

This  molecular  lowering  is  calculated  from  experiments  made 
with  more  dilute  solutions  in  the  following  manner.  Suppose  a 
solution  of  cane-sugar  (C12H22On,  mol.  wt.  =  342)  in  water,  con- 
taining 2*5  grams  in  100  grams  of  water,  is  found  by  experiment  to 
freeze  at  -o'i37°.  Then  — 

2*5  :  342  !  !  o'i37  =  187  =  molecular  depression 
or,  molecular  depression  =  — 

o 

*  This  does  not  hold  true  in  the  case  of  substances  whose  molecules  either 
dissociate  or  undergo  some  molecular  aggregation  in  the  particular  solvent. 
Thus,  in  the  case  of  water  as  solvent,  it  is  found  that  such  things  as  strong  acids 
and  bases  and  their  salts—  electrolytes,  in  other  words—  produce  a  molecular 
depression  much  greater  than  18-5.  This  will  be  seen  from  the  following 
table  (Raoult):— 


Non-electrolytes. 

Glycerol 

Mannite 

Sugar 

Ethyl  alcoho 

Pyrogallol 

Acetamide 


In  the  case  of  non-electrolytes  the  phenomenon  is  concerned  only  with 
the  molecules  of  the  substance  dissolved  in  water,  while  with  electrolytes  a 
certain  proportion  of  the  molecules  (depending  upon  the  degree  of  concentra- 
tion) are  believed  to  be  dissociated  into  their  ions. 


Mol. 

Mol. 

depression. 

Electrolytes. 

depression. 

l8'5 

Sulphuric  acid 

38*2 

18*1 

Hydrochloric  acid 

39'i 

...     18-6 

Nitric  acid  

35'3 

...     16-6 

Sodium  hydroxide 

...     i8'o 

Sodium  chloride    ... 

35'1 

...     177 

Potassium  chloride 

33-8 

444  Physico  chemical  Determinations. 

where  m  —  molecular  weight  of  the  dissolved  substance, 

/  =  observed  depression  of  the  freezing-point  of  the  solvent, 

and 

g  =  grams  of  substance  in  100  grams  of  the  solvent. 
Having  once  established  the  constant  representing  the  molecular 
lowering  of  freezing-point  for  any  particular  solvent,  we  are  then  in 
possession  of  all  the  data  for  determining  the  molecular  weight  of 
a  substance  which  will  dissolve  in  that  liquid  ;  for,  substituting  C 
for  the  constant  in  the  above  formula,  we  get — 

Or 


The  determination  may  be  made  by  means  of  the  apparatus 
shown  in  Fig.  101.  This  consists  of  a  delicate  thermometer 
graduated  into  T§oth*  °f  degrees,  which  is  fitted  by  means  of  a  cork 
into  the  mouth  of  a  large  test-tube,  A.  The  same  cork  also  carries 
a  short  piece  of  narrow  glass  tube,  C,  through  which  the  wire 
stirrer  passes.  The  tube  A  fits  into  a  wider  tube  (a  boiling-tube), 
B,  by  means  of  a  cork  or  a  short  piece  of  wide  rubber  tube.  The 
whole  is  supported  by  a  clamp  and  retort  stand,  and  can  be  lowered 
into  or  raised  out  of  the  beaker  or  glass  jar,  F,  containing  the 
freezing  mixture.* 

As  a  convenient  exercise  in  the  use  of  the  apparatus,  the 
molecular  weight  of  sugar  may  be  determined  by  the  lowering  of 
the  freezing-point  of  water. f 

Adjusting  the  Thermometer. — The  thermometer  employed 
differs  from  ordinary  thermometers,  in  that  it  has  an  adjustable 
zero.  The  scale  covers  a  range  of  usually  about  7  or  8  degrees, 
and  in  the  most  recent  forms  the  zero  is  at  the  upper  end  of  the 
scale.  The  total  quantity  of  mercury  in  the  thermometer  is  usually 
greater  than  is  required  to  bring  the  column  to  zero  when  the  bulb 
is  placed  in  ice  ;  hence  it  is  necessary  to  so  get  rid  of  some  of  the 
mercury  that  the  freezing-point  of  the  solvent  to  be  employed  (in 
the  present  example,  water)  shall  fall  somewhere  near  the  top  of 
the  scale— not  necessarily  exactly  at  zero,  but  conveniently  near  it. 

*  The  Beckmann  apparatus  usually  figured  in  instrument  makers'  price 
lists,  contains  a  number  of  unnecessary  appendages  which  give  it  an  appear- 
ance of  complication,  and  considerably  add  to  its  cost.  It  is  only  necessary  to 
purchase  the  thermometer  ;  for  the  rest,  two  boiling-tubes  and  a  beaker  or 
glass  jar  are  all  that  is  required. 

f  When  the  solvent  employed  is  hygroscopic  (glacial  acetic  acid,  for 
example),  the  glass  tube  c,  through  which  the  stirrer  passes,  should  be  replaced 
by  a  f-  tube ;  and  a  slow  stream  of  air,  dried  by  means  of  calcium  chloride  or 
sulphuric  acid,  allowed  to  pass  in  through  the  side  branch  and  escape  at  the  top. 
In  this  way  the  air  which  is  introduced  by  the  movement  of  the  stirrer  is  dry  air. 


Molecular   Weight  by  Freezing-point  Method      445 


FIG.  101. 


2   G 


446  Physico-chemical  Determinations. 

To  accomplish  this,  the  bulb  is  placed  in  water  3  or  4  degrees 
warmer  than  the  freezing-point  of  the  solvent,  that  is  to  say,  about 
+  3°,  since  water  is  to  be  the  solvent.*  The  effect  of  this  is  to 
send  the  mercury  up  until  it  begins  to  enter  the  wider  part  of  the 
bent  stem  E.  By  gently  tapping  the  side  of  the  instrument  with 
the  finger-tips,  a  small  piece  of  the  mercury  becomes  detached,  and 
falls  to  the  bottom  of  E,  so  that  when  the  bulb  is  then  placed  in 
ice,  the  top  of  the  column  of  mercury  will  come  to  rest  upon  the 
scale  somewhere  near  the  top  graduations. 

Into  the  tube  A,  perfectly  clean  and  dry,  a  weighed  quantity  of 
water  (about  20  to  25  grams)  is  introduced — either  by  weighing 
the  tube,  first  empty,  and  then  with  the  water,  or  by  weighing  the 
liquid  out  from  a  small  flask.  When  the  cork  carrying  the  ther- 
mometer and  stirrer  is  inserted,  the  bulb  should  be  entirely 
immersed  in  the  water.  Tube  A  is  then  placed  inside  tube  B,  and 
the  whole  system  lowered  into  a  freezing  mixture  of  ice  and  salt  in 
the  jar  F.  The  object  of  the  outer  tube  B  is  to  render  the  cooling 
more  gradual  and  uniform  than  would  be  the  case  if  it  were  directly 
dipped  into  the  freezing  mixture.f  As  the  water  cools,  it  should  be 
kept  gently  moving  by  means  of  the  stirrer,  the  handle  of  which 
should  be  thrust  into  a  small  cork  to  prevent  the  warmth  of  the 
hand  from  being  conducted  to  the  liquid.  The  temperature  will 
fall  considerably  below  the  freezing-point  before  solidification 
actually  sets  in,  but  the  moment  the  supercooled  liquid  begins  to 
freeze,  the  mercury  in  the  thermometer  will  run  up,  and  finally 
come  to  rest.  This  point,  which  may  be  -conveniently  read  by 
means  of  a  lens,  is  taken  as  the  freezing-point  of  the  water.  The 
observation  should  be  repeated  two  or  three  times,  by  removing 
tube  A,  and  gently  warming  it  with  the  hand  until  the  ice  is  re- 
melted,  and  then  going  through  the  freezing  process  again. 

When  the  exact  freezing-point  of  the  solvent  is  established,  a 
weighed  quantity  of  pure  cane-sugar  is  introduced  by  lifting  the 
cork  with  the  thermometer,  and  tipping  the  powdered  substance 
out  of  a  weighing-bottle. \  When  the  substance  has  completely 

*  If  benzene  (M.P.  6°)  were  to  be  used  as  the  solvent,  then  about  3°  above 
its  freezing-point ;  while  if  acetic  acid  (M.  P.  i6'5°),  then  a  temperature  of  about 
20°  would  be  suitable. 

t  It  is  sometimes  convenient,  and  a  saving  of  time,  to  cool  the  water  in  A 
down  to  the  freezing-point  by  first  dipping  the  tube  direct  into  the  freezing 
mixture.  Then  removing  it,  wiping  it  dry,  and,  when  the  contents  have  again 
entirely  remelted,  placing  it  inside  tube  B,  and  then  lowering  the  whole  into 
the  mixture. 

\  In  the  form  of  apparatus  usually  figured,  tube  A  has  a  branch  tube  blown 
into  it  near  the  top,  which  is  intended  for  the  introduction  of  the  substance ; 


Molecular    Weight  by  Freezing-point  Method.    447 

dissolved  (the  process  being  aided  by  means  of  the  stirrer),  the 
system  is  lowered  into  the  freezing  mixture,  and  the  freezing-point 
of  the  solution  determined  as  before,  by  two  or  three  repetitions  of 
the  operation. 

For  example  : — 

Weight  of  water  taken 20' 1 15  grams 

Freezing-point  indicated  by  thermometer  —  r  1 1 5° 

Weight  of  sugar  introduced       1*020  gram 

Freezing-point  of  solution         —  I"395° 

2O'I  I  C   X    IOO 

Percentage  strength  of  solution  =  -   — ± — —     =  5 '07  —  g 

lowering  of  freezing-point  =  1*395  —  ni5  =  o'28  =  / 
molecular  lowering  for  water  =  i8'5  —  c 

therefore  molecular  weight  =   -        ~ —  =  334*9 

O'2O 

When  a  supercooled  solution  begins  to  freeze,  and  the  temperature  runs  up 
to  a  point  at  which  the  thermometer  remains  for  a  time  practically  stationary, 
it  will  be  seen  that  a  certain  quantity  of  the  solvent  has  crystallised  out.  With 
the  same  solution,  the  amount  which  so  separates  depends  upon  the  extent  to 
which  it  was  previously  supercooled.  This  separation  of  a  portion  of  the 
solvent  obviously  renders  the  remaining  liquid  a  little  more  concentrated  than 
the  original  solution,  and  from  this  it  follows  (i)  that  the  observed  temperature 
is  slightly  lower  than  the  true  freezing-point  of  a  solution  having  the  original 
strength  ;  and  (2)  that,  unless  the  extent  to  which  the  liquid  is  supercooled  is  the 
same  each  time  the  operation  is  repeated,  the  results  obtained  will  not  exactly 
agree.  To  obviate  this,  it  is  necessary  to  make  a  correction  for  the  amount  of 
solvent  which  thus  separates. 

Let  t=  number  of  degrees  to  which  the  solution  is  supercooled, 

S  =  specific  heat  of  the  solution, 

L  —  latent  heat  of  the  solvent, 

M  =  mass  of  the  solution,  and 

m  —  mass  of  solvent  which  separates  out ; 

Then  S  x  t  x  M  =  Lm 
and  therefore  jrr(the  fraction  which  separates)  =  p 

For  practical  purposes,  the  specific  heat  of  the  solution  may  be  taken  as  the 
same  as  that  of  the  solvent ;  that  is  to  say,  the  specific  heat  of  the  small 
quantity  of  the  dissolved  substance  may  be  neglected.  Therefore,  in  the  case  of 

S#  / 

water,  S  —  i  and  L  =  80  ;    hence  -y-  becomes  =-.     That  is  to  say,   for  each 

degree  Centigrade  to  which  the  solution  is  supercooled,  &  of  the  solvent 
separates  out  when  the  freezing-point  is  indicated,  and  the  solution  thereby 
becomes  more  concentrated  to  the  same  extent.  And  since  the  depression  of 
the  freezing-point  is  proportional  to  the  concentration,  the  observed  depression 
will  be  too  low  by  the  same  fraction. 

but,  as  a  matter  of  fact,  the  substance  is  more  conveniently  introduced  at  the 
top  as  here  described,  and  with  the  greater  certainty  that  the  whole  of  it  actually 
falls  into  the  liquid. 


448  Physico-chemical  Determinations. 

Applying  this  correction  in  the  example  given  above  — 

The  lowest  temperature  to  which  the  mercury  sank  =  —2-695 
observed  freezing-point  =  —  i  -395 

therefore  degrees  of  supercooling  =       1-3 
Hence  -j-  becomes  —  *  =  o  '01625 

therefore  o'28  x  0*01625  =  '°°45S     • 

and  o'28  —  '00455  —  °'27545  =  corrected  depression 
Substituting  this  new  value  for  /  in  the  first  calculation  gives  — 


=  34J'3  -  corrected  molecular  weight. 


. 

7.  Determination  of  Molecular  Weight  by  the  "  Boil- 
ing-point Method."  —  Just  as  the  freezing-point  of  a  liquid  is 
lowered  by  the  presence  of  dissolved  substances,  so  the  boiling- 
point  is  raised  from  the  same  cause,  provided  the  substance  is 
either  non-volatile  or  exerts  no  appreciable  vapour-pressure  at  the 
boiling-point  of  the  solution. 

The  elevation  of  the  boiling-point  of  a  liquid  caused  by  the  solu- 
tion in  100  grams  of  it  of  a  gramme-molecule  of  such  a  substance 
is  termed  the  molecular  elevation,  or  the  molecular  rise  of  the 
boiling-point  ;  and  (with  the  same  reservations  as  to  substances 
which  dissociate,  as  applied  in  the  freezing-point  method)  for  one 
and  the  same  liquid  this  value  is  practically  a  constant. 

In  the  following  table  are  given  the  boiling-points  and  the  mole- 
cular rise  of  boiling-point  of  a  few  common  solvents  :  — 

Mol.  rise  of 
Boiling-point.  boiling-point. 

Water  .........   1000°     ......       5-2° 

Benzene         .........     80-5°     ......     27-0° 

Chloroform  .........     6i'2°     ......     36-6° 

Carbon  disulphide  ......     46*2°     ......     237° 

Ether  .........     34-9°     ......     2ri° 

The  operation  may  be  carried  out  by  means  of  the  "  Beck- 
mann"  apparatus,  shown  in  a  dissected  condition  in  Fig.  102. 

A  weighed  quantity  of  the  solvent  is  introduced  into  tube  A  — 
in  the  case  of  very  volatile  liquids,  such  as  ether,  by  weighing  the 
tube  with  the  ether,  the  tube  being  closed  by  corks.  This  tube  is 
similar  to  tube  A  in  the  freezing-point  apparatus,  except  that  here 
the  branch  tube  is  necessary  in  order  to  attach  a  small  reflux  con- 
denser, Ci,  to  return  the  liquid  which  would  otherwise  escape  as 
vapour  during  the  boiling.  Small  glass  beads,  b,  are  placed  in  the 
tube  to  a  depth  of  about  20  mm.  (a  short  inch),  and  the  quantity  of 
liquid  employed  must  be  such  that  the  bulb  of  the  thermometer  is 


Molecular  Weight  by  Boiling-point  Method.     449 


well  covered  when  the  bottom  of  it  nearly  touches  the  beads. 
Sometimes  a  short  piece  of  platinum  is  fused  through  the  bottom  of 
tube  A,  as  seen  at  /,  with  a  view  to  aid  in  the  heating. 

The  thermometer  is  similar  to  that  employed  in  a  freezing-point 
determination,  and  must  be  previously  adjusted,  so  that  the  boiling- 


FIG.  U2. 

point  of  the  liquid  to  be  used  may  be  indicated  upon  the  lower 
parts  of  the  graduated  scale.  The  zero  in  this  case  is  at  the  bottom 
of  the  scale. 

Tube  A  is  then  placed  within  the  jacket-tube  B,  which  is  open 
at  both  ends,  but  having  a  "jacket"  blown  upon  it,  in  which  can 
be  boiled  a  small  quantity  of  the  same  liquid  as  is  in  lube  A. 


45°  Physico  chemical  Determinations. 

This  jacket  is  also  connected  to  a  reflux  condenser,  in  order  to 
return  the  vaporised  liquid.  The  object  of  this  tube  is  to  surround 
tube  A  with  an  envelope  which  shall  have  a  temperature  as  nearly 
as  possible  the  same  as  that  within  A  itself. 

The  apparatus  is  heated  upon  the  asbestos  support  D.  This 
somewhat  resembles  the  lid  of  a  box,  provided  with  two  asbestos 
chimneys  for  carrying  away  the  waste  heat  from  the  lamps.  From 
the  centre  hole,  over  which  tube-B  is  placed,  there  depends  a  short 
wide  asbestos  tube,  and  the  lamp  (or  lamps)  are  placed  so  that  their 
small  flames  are  outside  this  tube,  and  therefore  not  directly  below 
the  glass  tubes.  The  heat  must  be  so  regulated  that  a  steady  and 
brisk  boiling  of  the  liquid  shall  take  place  ;  any  irregularity  or 
"  bumping  "  is  quite  fatal  to  the  success  of  the  operation.  It  is 
desirable  to  still  further  shield  the  apparatus  by  a  screen  of  asbes- 
tos paper.  When  the  mercury  is  stationary  (after  one  or  two 
gentle  taps  upon  the  stem  of  the  thermometer  with  the  finger-tips), 
the  reading  is  taken  —  using  a  pocket-lens  —  after  which  the  appa- 
ratus is  allowed  to  cool,  and  the  operation  repeated,  if  necessary, 
once  or  twice. 

When  the  exact  boiling-point  of  the  liquid,  as  registered  by  the 
thermometer,  has  been  ascertained,  a  weighed  quantity  of  the  sub- 
stance whose  molecular  weight  is  to  be  determined  is  then  intro- 
duced, and  after  it  has  entirely  dissolved,  the  boiling-point  of  the 
liquid  is  again  taken.  The  molecular  weight  is  calculated  from 
the  formula  (similar  to  that  for  freezing-points)  — 


where   C  =  constant  (or  mol.  elevation  of  boiling-point  of  solvent), 
g  =  grams  of  substance  in  100  grams  of  solvent,  and 
R  =  observed  rise  of  boiling-point. 

The  method  is  less  capable  of  yielding  accurate  results  than  the 
freezing-point  method,  partly  owing  to  conditions  of  experiment, 
and  partly  to  variations  of  atmospheric  pressure;  but  since  the 
object  is  usually  merely  to  decide  between  two  values,  one  of  which 
is  a  multiple  of  the  other,  a  considerable  experimental  error  is  not 
of  serious  consequence. 

As  an  exercise  for  practice  in  the  method,  the  molecular  weight 
of  turpentine  (C10H16  =  136)  may  be  determined,  using  ether  as 
the  solvent.* 

*  Water  is  unfortunately  not  a  convenient  solvent,  partly  owing  to  disso- 
ciation suffered  by  so  many  substances  in  solution  in  this  liquid,  and  partly  on 


Molecular   Weight  by  Boiling-point  Method.       451 

EXAMPLE  :  — 

Weight  of  ether  taken    ............     18-45  grams 

Boiling-point  indicated  by  the  thermometer  ...       3-25° 
By  means  of  a  bent  pipette,  a  weighed  quan- 
tity  of  turpentine  was    introduced   at    the 
branch  tube      ...............       0-8650  gm. 

Boiling-point  of  the  solution     ...  ...       3  94.° 

0-86SO  x  100 

Percentage  strength  of  solution  =  -  •—  -    =  4'688  =  g 

' 


rise  of  boiling-point  =  3*94  —  3*35  =  0*69  =  R 
mol.  rise  of  boiling-point  for  ether  =  21*1  =  C 

Hence  mol.  wt.  of  turpentine  =  -  -  -  =  143*3 

0-69 

account  of  its  very  low  molecular  elevation  of  boiling-point.  A  volatile  liquid 
like  ether,  on  the  other  hand,  necessitates  very  careful  condensation.  The 
most  effective  condensers  are  Cribb's  double-surface  condensers  ;  and  the  water 
with  which  they  are  supplied  should  be  cooled  by  being  passed  through  a  few 
turns  of  "  compo  "  pipe  placed  in  a  bowl  of  ice.  The  temperature  of  the  water 
issuing  from  the  condenser  may  thus  easily  be  maintained  at  5°  or  6°  C. 


APPENDIX 

TABLE  OF  ATOMIC  WEIGHTS. 

Atomic  weights. 

2    "a        i"  i 


Name. 

.yjd 

8s  a 

iii 

stsss 

•S  S-5 

S  £  rt 

Name. 

°^2 

IJB 

H 

iss 

kg-3 

%*«> 

More 
exact 
values. 

Aluminium 

Al 

27 

27-04 

Molybden  u  m 

Mo 

96 

95  '9 

Antimony 

Sb 

120 

II9-6 

Nickel   

Ni 

59 

58-6 

Arsenic   ... 

As 

75 

74'9 

Niobium 

Nb 

937 

— 

Barium    ... 

Ba 

137 

I36-86 

Nitrogen             .. 

N 

14 

14*01 

Sery  Ilium 

Be 

9 

9-08 

Osmium 

Os 

191 

— 

Bismuth  ... 

Bi 

2075 

— 

Oxygen  ... 

0 

16 

15-96 

Boron 

B 

11 

10'9 

Palladium 

Pd 

106 

106-2 

Bromine  ... 

Br 

80 

79-76 

Phosphorus 

P 

31 

30-96 

Cadmium 

Cd 

112 

111-7 

Platinum 

Pt 

195 

!94'3 

Casium   ... 
Calcium  ... 

Cs 

Ca 

133 
40 

132-7 
39V 

Potassium    (ka-  \ 
Hum}  ...          / 

K 

39 

39-03 

Carbon    ... 

C 

12 

11-97 

Rhodium 

Rh 

10  i 

fc>4'i 

Cerium    ... 

Ce 

1412 

— 

Rubidium 

Rb 

85 

85-2 

Chlorine... 

Cl 

355 

35-37 

Ruthenium 

Ru 

103-5 

Chromium 

Cr 

52 

52"45 

Samarium 

Sm 

150 



Cobalt     

Co 

59 

58-6 

Scandium 

Sc 

44 

43'97 

Copper  (cuprum) 

Cu 

63 

63-18 

Selenium 

Se 

79 

78-87 

Didymium 

Di 

145 

-  — 

Silicon  ... 

Si 

28 

28-3 

Erbium  ... 

Er 

166 

— 

Silver     

Ag 

108 

107-66 

Fluorine... 

F 

19 

19-06 

Sodium  ... 

Na 

23 

22-995 

Gallium  ... 

Ga 

70 

69-86 

Strontium 

Sr 

87-3 

Germanium 

Ge 

72 

— 

Sulphur 

S 

32 

31-98 

Gold  (aunun}     ... 

Au 

197 

196-8 

Tantalum 

Ta 

182 

Hydrogen 

H 

1 

— 

Tellurium 

Te 

125 



Indium   ... 

In 

113 

H3'4 

Thallium 

Tl 

2037 



Iodine 

I 

127 

126-54 

Thorium 

Th 

232 



Iridium    ... 

Fr 

1925 

Tin        

Sn 

118 

ii7'35 

Iron  (ferrum)    ... 

Fe 

56 

55-88 

Titanium 

Ti 

48 

Lanthanum 

La 

138-5 

Tungsten 

W 

184 



Lead  (plumbum} 

Pb 

207 

206-39 

Uranium 

U 

239-8 



Lithium  ... 

Li 

7 

7-01 

Vanadium 

V 

51-1 



Magnesium 

Mg 

24 

23-94 

Vtterbium 

Yb 

173 



Manganese 

Mn 

55 

54-8 

Yttrium 

Y 

896 



Mercury  (hydrar-\ 
gyrum               } 

Hg 

200 

199-8 

Zinc       
Zirconium 

Zn 
Zr 

65 
904 

64-88 

454  Reagents. 


REAGENTS. 

It  is  important  that  the  reagents  used  in  qualitative  analysis  should 
be  made  up  of  known  definite  strength,  and  that  the  strength  should  be 
indicated  upon  the  bottles.*  It  is  a  great  additional  advantage,  also,  that 
the  reagents  should,  as  far  as  possible,  have  either  a  uniform  strength 
(that  is  to'say,  equal  volumes  should  possess  equal  chemical  values,  or  be 
chemically  equivalent),  or,  when  this  is  not  possible  or  desirable,  the 
strength  should  bear  some  simple  ratio  to  the  equivalent  strength. 

By  this  is  not  meant  that  the  reagents  are  to  be  prepared  with  an 
exactness  in  any  way  comparable  to  that  required  in  making  up  the 
standard  solutions  used  in  volumetric  analysis.  For  qualitative  analysis, 
it  is  enough  that  the  reagents  bear  a  rough  approximation  to  the  exact 
strength  indicated. 

Unless  some  of  the  reagents  were  to  be  inconveniently  dilute,  it  would 
not  be  possible  to  have  them  all  chemically  equivalent,  owing  to  differences 
of  solubility.  For  example,  a  solution  of  oxalic  acid,  which  is  saturated 
at  IO°,  is  only  equivalent  to  sulphuric  acid  consisting  of  I  volume  of  strong 
acid  in  about  18  volumes  of  water. 

In  volumetric  analysis,  standard  solutions  are  designated  normal 
solutions  when  they  are  of  such  a  strength  that  I  litre  contains  a  weight  of 
the  reagent  in  grams,  equal  to  the  chemical  equivalent  of  that  reagent. 
Thus,  normal  sulphuric  acid  contains  49  grams  of  IJ2SO4  per  litre  ;  normal 
hydrochloric  acid  36*5  grams  of  HCl  per  litre,  and  so  on  (see  p.  313). 
The  signature  employed  to  denote  fhese  normal  or  equivalent  solutions 
is  the  letter  N. 

For  qualitative  analysis  a  convenient  strength  for  the  common  acids  and 
alkalies  and  general  reagents  is  five  times  this  normal  strength.  These 
are  distinguished  by  the  signature  5N. 

A  large  number  of  other  reagents,  such  as  are  used  for  special  tests,  are 
conveniently  prepared  of  normal  strength,  and  should  bear  the  signature  N  ; 
while  others,  which  for  special  reasons  (such  as  slight  solubility,  e.g.  calcium 
sulphate;  or  the  delicacy  of  their  reactions,  e.g.  ammonium  thiocyanate) 
should  be  more  dilute,  are  prepared  of  suitable  strengths,  given  in  the 
following  list  of  reagents,  and  the  signature  of  the  strength  in  all  cases 
should  be  indicated  upon  the  labels. 

Thus,  in  the  case  of  calcium  sulphate,  a  saturated  solution  is  employed  ; 
this  contains  one-thirtieth  of  the  quantity  of  calcium  sulphate  required  to  pro- 
duce a  normal  or  equivalent  solution,  and  its  strength  is  denoted  by  - . 

*  It  is  greatly  to  be  regretted  that  in  many  laboratories  the  practice  still 
exists  of  providing  the  student  with  reagents  concerning  whose  strength  he  is 
left  absolutely  in  the  dark.  The  effect  of  this  upon  the  analytical  work  is  only 
too  familiar  to  those  who  have  had  much  practical  experience  in  teaching. 


Reagents  of  Standard  Strength.  455 

Again,  ammonium  thiocyanate  being  so  delicate  a  reagent  for  the 
detection  of  ferric  salts,  a  solution  of  one-fifth  of  the  equivalent  strength 
is  conveniently  strong.  This  is  indicated  upon  the  label  by  the 

N* 
signature  —  . 

Reagents  of  Five  Times  Normal  Strength.     Signature,  5N. 

Sulphuric  acid,  H2SO4.  Equivalent  =  49.  Ordinary  concentrated 
acid,  sp.  gr.  1*84,  has  a  strength  36  times  normal,  and  may  be  designated 
36N.  One  volume  of  strong  acid  diluted  with  6  volumes  of  water  =  5N. 

Nitric  acid,  HNO3.t  Equivalent  =  63.  Concentrated  acid,  sp.  gr. 
i '42  =  i6N.  Five  volumes  of  strong  acid  diluted  with  n  volumes  of 
water  =  5N. 

Hydrochloric  acid,  HCl.f  Equivalent  =  36-5.  Concentrated  acid, 
sp.  gr.  1*16  =  loN.  One  volume  of  strong  acid  diluted  with  I  volume  of 
water  =  5N. 

Acetic  acid,  C2H4O2.  Equivalent  =  60.  Glacial  acid,  M.P.  10°  C., 
=  i;N.  One  volume  of  glacial  acid  diluted  with  2\  volumes  of 
water  =  5N. 

Potassium  hydroxide,  KHO.     Equivalent  =  56. 

280  grams  dissolved  in  water  and  diluted  to  one  litre  =  5N. 

Sodium  hydroxide,  NaHO.     Equivalent  =  40. 

200  grams,  dissolved  in  water  and  diluted  to  one  litre  =  sN. 

Ammonia  (solution  in  water,  NH4HO).     Equivalent  =  35. 

The  strong  solution,  sp.  gr.  o'88o  —  2oN.  One  volume  of  the  strong 
solution  diluted  with  3  volumes  of  water  =  5N. 

Ammonium  sulphide,  (NH4)2S.    Equivalent  =  34. 

600  c.c.  of  5N  ammonia  are  saturated  with  sulphuretted  hydrogen; 
this  gives  hydrogen  ammonium  sulphide  H(NH4)S.  This  is  made  up  to 
i  litre  by  adding  5N  ammonia.  (This  reagent  is  slowly  decomposed  by 
atmospheric  oxygen,  ammonia  is  evolved,  and  yellow  ammonium  sulphide 
is  formed  (NH4)2S2.) 

Sodium  sulphide,  Na2S.     Equivalent  =  39. 

200  grams  of  sodium  hydroxide  are  dissolved  in  800  c.c.  of  water. 

400  c.c.  of  this  solution  are  saturated  with  sulphuretted  hydrogen,  and 
the  remaining  half  added,  together  with  water  sufficient  to  make  the 

*  The  system  of  standard  reagents  for  use  in  qualitative  analysis,  upon 
which  the  instructions  here  given  are  based,  was  devised  by  Reddrop,  and  is 
published  in  full  detail  in  the  Chemical  News,  May,  1890.  Mr.  Reddrop 
employs  the  signature  E  (equivalent]  instead  of  N  (normal]  in  his  system  ;  but 
as  the  designation  normal  is  the  term  almost  universally  employed  for  such 
solutions  when  used  for  volumetric  analysis,  it  appears  more  harmonious,  and 
brings  qualitative  analysis  more  into  line  with  quantitative  methods,  to  adopt  a 
uniform  nomenclature  throughout. 

t  Aqua  regia  is  prepared  as  required  by  mixing  i  volume  of  i6N  nitric  acid 
with  3  volumes  of  loN  hydrochloric  acid. 


456  Reagents  of  Standard  Strength. 

volume  up  to  I  litre.  When  hydrogen  sodium  sulphide,  HNaS,  is 
required,  the  sodium  hydroxide  is  simply  saturated  with  sulphuretted 
hydrogen,  without  the  further  addition  of  sodium  hydroxide. 

Ammonium  chloride,  NH4C1.     Equivalent  =  53-5. 

267*5  grams  of  the  salt  dissolved  in  water  and  diluted  to  one  litre  =  5N. 

Ammonium  carbonate,*  (NH4)2CO3.     Equivalent  =  48. 

200  grams  of  ammonium  sesquicarbonate  (commercial  carbonate)  dis- 
solved in  350  c.c.  of  5N  ammonia,  and  diluted  with  water  to  one  litre 
=  SN. 

Ammonium  acetate,  (NH4)C2H3O2.    Equivalent  =  77. 

300  c.c.  of  glacial  acetic  acid  (i7N)  neutralised  with  strong  ammonia 
and  diluted  to  one  litre  =  5N. 

Reagents  of  Normal  Strength.    Signature  N. 

The  following  normal  reagents  are  prepared  by  dissolving  the  equi- 
valent weight  in  grams  of  the  various  salts  in  water,  and  diluting  to  one 
litre.  In  all  cases  the  nearest  whole  number  to  the  equivalent  weight 
may  be  taken  as  sufficiently  exact. 

Ammonium  sulphate,  (NH4)2SO4.  Equivalent  weight  ...     66*0 

Barium  chloride,  BaC!2,2H2O.                                  ,,  ,,  ...  122*0 

Calcium  chloride,  CaCl2,6H2O.                                „  ,,  ...  109-5 

Copper  sulphate,  CuSO4, 5 H2O.                                ,,  ,,  ..    12475 

Ferric  chloride,  FeCl3.                                                ,,  ,,  ...     54*17 

Ferrous  sulphate,  FeSO4,7H2O.                               ,,  ,,  ...  139*0 

Hydrogen disodium phosphate,  HNa2PO4,i2H2O.  ,,  ,,  ...  119*3 

Lead  acetate  Pb(C2H3O2)2,3H2O.                             ,,  ,,  ...189*5 

Lead  nitrate,  Pb(NO3)2.                                               ,,  ,,  ...165*5 

Magnesium  chloride,  MgCJ2,6H2O.                         ,,  ,,  ...  101*5 

Magnesium  sulphate,  MgSO4,7H2O.                        ,,  ,,  ...  123*0 
"Magnesia  mixture"  (MgCl2,  NH4C1,  and  NH4HO). 

68    grams    of    MgCl2,6H2O,    together    with    165  grams  of 

NH4C1,  are  dissolved  in   300  c.c.  of  water;   300  c.c.  of 

5N  ammonia  are  added,  and  the  solution  diluted  with  water 
up  to  one  litre. 

Potassium  chromate,  K2CrO4.  Equivalent  weight  ...     97*25 

Potassium  cyanide,  KCy.                                           ,,  ,,  ...     65  o 

Potassium  ferrocyanide,  K4FeCy6,3H2O.                ,,  ,,  ...  105*5 

Potassium  ferricyanide,  K3FeCy6.                             ,,  ,,  ...  109*7 

Potassium  sulphate,  K2SO4.                                       ,,  ,,  ...     87*0 

Sodium  acetate,  NaC2H3O2,3H2O.                          ,,  „  ...136-0 

*  Hydrogen  ammonium  carbonate,  H(NH4)CO3,  cannot  be  prepared 
stronger  than  3^  (see  below). 


Reagents.  457 

Sodium  acetate  and  acetic  acid. 

136  grams  of  the  salt  dissolved  in  800  c.c.  of  water  and  200 

c.c.  of  glacial  acetic  acid  added. 

Sodium  carbonate,*  Na2CO3,ioH2O.  Equivalent  weight  ...  143*0 

Sodium  thiosulphate,  Na2S2O3,5H2O.  Equivalent  weight  ...   124-0 

Stannous  chloride,  SnCl2,2H2O.  ,,  „       ...    112-5 

112  grams  of  the  salt  are  dissolved  in  200  c.c.  of  $N  HC1, 
and  the  solution  diluted  with  water  to  one  litre.  Fragments 
of  granulated  tin  should  be  placed  in  the  solution. 

Beagents  of  Various  Strengths. 

Ammonium  oxalate,  (NH4)2C2O4,2H2O.        Equivalent  weight  ...     80-0 

N 
40  grams  of  the  salt  in  one  litre  of  water  =  —  solution. 

Ammonium  bicarbonate,  H(NH4)CO3. 

Saturation  solution  =  3N   solution.      Prepared    by   passing 

excess  of  CO2  into  3N  ammonia. 
Ammonium  thiocyanate,  NH4CyS.  Equivalent  weight  ...     76-0 

15  grams  of  the  salt  in  one  litre  of  water  =  —  solution. 
Barium  hydroxide,  Ba(HO)2,8H2O.  Equivalent  weight  ...  157*5 

31  grams  in  one  litre  of  water  =  —  solution. 
Barium  nitrate,  Ba(NO3)2.  Equivalent  weight  ...  130-5 

65  grams  dissolved  in  one  litre  of  water  =  —  solution. 

Bromine  water,  Br.  Equivalent  weight  ...     80*0 

Saturated  solution,   obtained    by  shaking   up   an  excess  of 

N 
bromine  with  water.     Gives  a  solution  of —  strength. 

Calcium  hydroxide  (lime-water)  Ca(HO)2.      Equivalent  weight  ...     37*0 

A  saturated  solution,  obtained  by  shaking  up  an  excess  of 

lime    with    water  and    allowing    the    mixture    to    settle, 

=  about  —  solution. 
20 

Calcium  sulphate,  CaSO4.  Equivalent  weight  ...     66 'O 

A  saturated  solution,  obtained  by  shaking  up  an  excess  of 

calcium  sulphate  with  water,  =  —  solution. 
Chlorine  water,  Cl.  Equivalent  weight  ...     35*5 

N^ 

A  saturated  solution  =  — . 
5 

*  A  strong  solution  of  sodium  carbonate,  sN  strength,  should  also  be 
prepared  by  dissolving  429  grams  of  the  salt  in  the  litre.  This  is  practically 
a  saturated  solution  at  ordinary  temperatures. 


458  Reagents. 

Mercuric  cl.loride,  HgCl2.  Equivalent  weight  ...   135-5 

XT 

27  grams  dissolved  in  one  litre  of  water  = —  solution. 

Mercurous  nitrate,  Hg2(NO3)2,2H2O.  Equivalent  weight  ...  280*0 

56  grams  dissolved  in  40  c.c.  of  5N  HNO3  and  diluted  to  one 

N 
litre  =    -  solution.     A  small  quantity  of  mercury  should  be 

placed  in  the  bottle. 
Potassium  iodide,  KI.  Equivalent  weight  ...  166  o 

N 
33  grams  dissolved  in  one  litre  of  water  =       solution. 

Potassium  metantimoniate,  KSbO3,3H2O.      Equivalent  weight  ...  291-0 
A  saturated  solution,  obtained  by  shaking  up  an  excess  of  the 

N 
salt  with  water,  =  -^ 

Silver  nitrate,  AgNO3.  Equivalent  weight  ...  170*0 

3'4  grams  dissolved  in  100  c.c.  =—  solution. 

Fusion  Mixture. — Pure  anhydrous  sodium  and  potassium  carbonates 
are  intimately  mixed  together  in  the  proportion  of  their  equivalent 
weights.  Sodium  carbonate  53  parts,  and  potassium  carbonate  69  parts. 


Appendix. 


459 


TABLE  OF  HARDNESS. 
Degrees  of  hardness  expressed  as  parts  of  CaCO3  per  100,000. 


C.C.  of 
soap 
solution. 

CaCO3 
per  100,000 
parts. 

C.C.  of 
soap 
solution. 

CaCOa 
per  100,000 
parts. 

C.C.  of 
soap 
solution. 

CaCOg 

per  100,000 
parts. 

C.C.  of 
soap 
solution. 

CaCOa 
per 

100,000 

parts. 

I'O 

•48 

5*3 

6'oo 

90 

1  1  'SO 

I3-0 

1  8  -02 

•I 

•63 

•14 

'I 

•95 

•I 

•17 

•2 

79 

•2 

•29 

'2 

I2"II 

*2 

'33 

•3 

'95 

'3 

'43 

'3 

•26 

'3 

•49 

'4 

I'll 

'4 

'57 

'4 

•41 

"4 

•65 

I 

•27 
'43 

1 

'1 

•56 
•71 

I 

•81 
•97 

7 

•56 

7 

7-00 

7 

•86 

7 

19-13 

•8 

•69 

•8 

•14 

•8 

13-01 

•8 

•29 

"9 

•82 

'9 

•29 

'9 

•16 

'9 

'44 

2'0 

'95 

6-0 

'43 

IO'O 

•31 

14-0 

•60 

•I 

2-08 

•i 

'57 

•i 

•46 

•i 

•76 

•2 

•21 

•2 

71 

'2 

•61 

'2 

•92 

"3 

'34 

'3 

•86 

'3 

•76 

'3 

20-08 

"4 

'47 

'4 

8-00 

'4 

•91 

'4 

•24 

'5 

•60 

•5 

'*4 

'5 

14-06 

'5 

•40 

•6 

73 

•6 

•29 

•6 

•21 

•6 

•56 

7 

•86 

7 

*43 

7 

'37 

7 

71 

•8 

•99 

•8 

'57 

•8 

•52 

•8 

•87 

'9 

3-12 

'9 

•71 

'9 

•68 

'9 

21-03 

3-0 

•25 

7-0 

•86 

I  I'O 

•84 

15-0 

•19 

•i 

•38 

•i 

9-00 

•I 

15-00 

•i 

•35 

"2 
'3 

•11 

•2 
'3 

•14 

•29 

'2 
'3 

•16 
•32 

'2 
•3 

$ 

'4 

77 

'4 

'43 

'4 

•48 

'4 

'85 

'5 

•90 

"57 

'5 

•63 

'5 

22-02 

•6 

•6 

71 

•6 

79 

•6 

•18 

7 

•16 

7 

•86 

7 

'95 

7 

'35 

•8 

•29 

•8 

lO'OO 

•8 

16-11 

•8 

•52 

'9 

'43 

'9 

•15 

'9 

•27 

'9 

•69 

4-0 

•57 

8-0 

•30 

I2'0 

'43 

16-0 

•86 

•i 

•71 

•i 

•45 

'I 

'59 

•2 

•86 

'2 

•60 

'2 

75 

'3 

5-00 

"3 

75 

•3 

•90 

'4 

'H 

'4 

•90 

'4 

17-06 

5 

•29 
•43 

:| 

11-05 

•20 

•6 

•22 

•38 

7 

'57 

7 

•35 

7 

'54 

•8 

•71 

•8 

•50 

•8 

70 

•9 

•86 

•9 

•65 

'9 

•86 

46o 


Appendix. 


EXPANSION  OF  WATER  BETWEEN  o°  AND  25°  C. 


Temperature 

Volume. 
(Vol.  at  4°  =  i  ) 

Density. 

(Density  at  4°  =  i.) 

0 

O 

I-OOOI2 

0-99988 

2 

1-00003 

0-99997 

4 

I'OOOOO 

I  00000 

6 

1-00003 

0-99997 

8 

I  -0001  I 

0-99989 

10 

1-00025 

Q'99975 

12 

1-00044 

0-99956 

13 

1-00055 

0-99945 

«4 

1-00068 

0-99932 

15 

1-00082 

0-99918 

15'S 

1-00089 

0-99911 

16 

1-00097 

0-99903 

17 

1-00113 

0-99887 

18 

1-00131 

0-99869 

19 

1-00149 

0-99851 

20 

1-00169 

0-99831 

22 

I-QO2I2 

0-99787 

24 

I  -00259 

0-99742 

25 

I  00284 

0-99717 

I 

The  coefficients  of  expansion  of  water  (the  values  in  the  middle  column) 
between  the  temperatures  o°  and  25°  may  be  calculated  from  the 
formula — 

V  =  V0  —  0-000061045^  +  0-000007  7 1 83/2  —  o-ooooooo3734/3 

where  V0  =  the  volume  at  o°.     In  the  above  table  V0  =  roooi2  ;  but  if 
the  coefficients  are  required  when  the  volume  at  o°  is  taken  as  unity, 


For  temperatures  above  25°  a  different  formula  is  required. 


Appendix. 


TENSION  OF  AQUEOUS  VAPOUR  IN  MILLIMETRES  OF  MERCURY. 
For  each  fifth  of  a  degree  from  5°  to  25°  C. 


o 

mm. 

0 

mm. 

0 

mm. 

0 

mm. 

5'° 

6°5 

I0'0 

9-2 

I5-0 

127 

2O  'O 

17-4 

•2 

6-6 

'2 

9'3 

•2 

I2'9 

'2 

'4 

6  7 

'4 

9'4 

•4 

I3-0 

"4 

I7'8 

•6 

6-8 

•6 

9'5 

•6 

I3-2 

•6 

18-0 

•8 

69 

•8 

97 

•8 

•8 

18-3 

6-0 

7-0 

II  0 

16-0 

I3-5 

21'0 

18-5 

'2 

7-1 

'2 

9-9 

'2 

137 

'2 

187 

*4 

7-2 

'4 

IO'I 

'4 

I3-9 

'4 

19*0 

•6 

7'3 

•6 

10-2 

•6 

I4-I 

•6 

19-2 

•8 

7  '4 

•8 

10-3 

•8 

I4-2 

•8 

19-4 

7-0 

7'5 

12  0 

io'5 

17-0 

14-4 

22  'O 

197 

'2 

7-6 

•2 

io'6 

•2 

•2 

19-9 

'4 

77 

'4 

107 

'4 

I4'8 

'4 

20'I 

•6 

•6 

10*9 

•6 

I5-0 

•6 

20-4 

•8 

7  9 

•8 

i  no 

•8 

I5-2 

•8 

20'6 

8-0 

8-0 

13-0 

1  1  '2 

18-0 

IS'4 

23-0 

20-9 

•2 

8-1 

•2 

11-3 

'2 

I5-6 

'2 

2I'I 

'4 

8-2 

'4 

11-5 

'4 

157 

*4 

2  1  '4 

•6 

8*3 

•6 

116 

•6 

•6 

2  17 

•8 

8'5 

•8 

11-8 

•8 

16-1 

•8 

21-9 

9-0 

8-6 

14-0 

11-9 

19-0 

16-3 

24-0 

22'2 

'2 

87 

'2 

'2 

166 

'2 

22  '5 

3 

8-8 
89 

'4 

•6 

I2'2 

'4 

•6 

16-8 
170 

6 

227 
23'0 

•8 

9-0 

•8 

I2'5 

•8 

17-2 

•8 

23'3 

25-0 

23'5 

Incases  where  the  tension  rises  0*1  mm.  for  a  rise  of  o'2°,  the  pressure 
for  the  intermediate  tenth  degree  may  be  taken  as  the  same  as  that  given 
for  the  temperature  immediately  preceding  it.  Thus,  for  the  temperature 
10-1°,  the  tension  9-2  mm.  will  be  taken.  For  very  accurate  work,  fuller 
tables  given  to  the  third  decimal  should  be  consulted. 


2  H 


462 


Appendix. 


FACTORS  FOR  REDUCING  GASEOUS  VOLUMES  TO  N.T.P. 

The  observed  volume,  when  multiplied  by  the  factor  corresponding  to 
the  temperature  and  the  pressure,  will  give  the  volume  reduced  to  o°  and 
760  mm.  See  footnote  on  p.  390. 


mm. 

9°- 

10°. 

n°. 

12°. 

13°. 

14°. 

720 

•9I7II 

•91388 

•91066 

•90746 

•90426 

•90113 

I 

•91839 

•91515 

•91192 

•90872 

•90552 

•90238 

2 

•91966 

•91642 

•91318 

•90998 

•90677 

•90363 

3 

•92093 

•91769 

•9H45 

•9II24 

•90803 

•90488 

4 

•92221 

•91896 

•9i57i 

•91250 

•90929 

•90614 

5 

•92348 

•92023 

•91698 

•91376 

•91054 

•90739 

6 

•92476 

•91150 

•91824 

•91502 

•Qii8o 

•90864 

7 

•92603 

•92277 

•9i95i 

•91628 

•91305 

•90989 

8 

•92730 

•92403 

•92077 

•91754 

•9143! 

•9III4 

9 

•92858 

•92530 

•92204 

•91880 

•91556 

•91239 

730 

•92985 

•92657 

•92330 

•92006 

•91682 

•91365 

i 

•93II2 

•92784 

•92457 

•92132 

•91808 

•91490 

2 

•93240 

•92911 

•92583 

•92258 

•91933 

•9I6I5 

3 

•93367 

•93038 

•92710 

•92384 

•92059 

•91740 

4 

•93495 

•93165 

•92836 

•92510 

•92184 

•91865 

5 

•93622 

•93292 

•92963 

•92636 

•92310 

•91990 

6 

'93749 

'93419 

•93089 

•92762 

•92436 

•92H5 

7 

•93X77 

•93546 

•93216 

•92888 

•92561 

•92241 

8 

•94004 

•93673 

•93342 

•93014 

•92687 

•92366 

9 

•94132 

•93800 

•93469 

'9W 

•92812 

•92491 

740 

•94259 

•93927 

'93595 

•93267 

•92938 

•92616 

•94386 

•94054 

•93722 

'93393 

•93064 

•92741 

2 

•945H 

•94180 

•93848 

•93519 

•93189 

•92866 

3 

•94641 

•94307 

'93975 

•93645 

•93315 

•92992 

4 

•94768 

•94434 

•94101 

•93771 

•93440 

'93"7 

5 

•94896 

•94561 

•94228 

•93897 

•93566 

•93242 

6 

•95023 

•94688 

'94354 

•94023 

•93692 

•93367 

7 

•95!5i 

•94815 

•94480 

•94149 

•93817 

•93492 

8 

•95278 

•94942 

•94607 

•94275 

•93943 

•93617 

9 

•95405 

•95069 

'94733 

•94401 

•94068 

•93742 

750 

'95533 

•95196 

•94860 

•94527 

•94194 

•93868 

i 

•95660 

•95323 

•94986 

•94653 

'943  i  9 

'93993 

2 

•95788 

•95450 

'95"3 

'94779 

'94445 

•94118 

3 

•95915 

•95577 

•95239 

•94905 

•94571 

•94243 

4 

•96042 

•95704 

•95366 

•95031 

•94696 

•94368 

5 

•96170 

•9583i 

•95492 

•95157 

•94822 

'94493 

6 

•96297 

•95958 

•95619 

•95283 

'94947 

•94619 

7 

•96424 

•96084 

'95745 

•95409 

•95073 

'94744 

8 

•96552 

•96211 

•95872 

'95535 

•95J99 

•94869 

9 

•96679 

•96338 

•95998 

•95661 

•95324 

'94994 

Factors  for  reducing  Gaseous  Volumes  to  N.T.P.     463 


mm. 

9° 

10°. 

n°. 

12°. 

IS0' 

14°. 

760 

•96806 

•96465 

•96125 

•95787 

•95450 

•95"9 

•96934 

•96592 

•96251 

•95913 

'95575 

•95244 

2 

•97061 

•96719 

•96378 

•96039 

•95701 

•95370 

3 

•97189 

•96846 

•96504 

•96165 

•95827 

'95495 

4 

•97316 

•96973 

•96631 

•96291 

•95952 

•95620 

5 

'97443 

•97100 

•96757 

•96417 

•96078 

'95745 

6 

•97571 

•97227 

•96884 

•96543 

•C6203 

•95870 

7 

•97698 

•97354 

•97010 

•96670 

•96329 

'95995 

8 

•97825 

•97481 

'    '97137 

•96796 

•96455 

•96120 

9 

'97953 

•97608 

•97263 

•96922 

•965^0 

•96246 

770 

•98080 

•97734 

•97390 

•97048 

•96706 

•96371 

•98208 

•97861 

•975i6 

•97174 

•96831 

•96496 

2 

•98335 

•979^8 

•97642 

•97300 

•96957 

•96621 

3 

•98462 

•98115 

•97709 

•97426 

•97083 

•96746 

4 

•98590 

•98242 

•97895 

•97552 

•97208 

•96871 

5 

•98717 

•98369 

•98022 

•97678 

'97334 

•96997 

6 

•98844 

•98496 

•98148 

•97804 

'97459 

•97122 

7 

•98972            -98623 

•98275 

•97930 

•97585 

•97247 

8 

•99099            -98750 

•98401 

•98056 

•97710 

•97372 

9 

•99227            -98877 

•98528 

•98l82 

•97836 

•97497 

780 

'99354     !     -99004 

•98654 

•98308 

•97962 

•97622 

mm. 

15°.        16°, 

i?°. 

18°. 

19°. 

20°. 

720 

•89800 

•89489 

•89180 

•88873 

•88569 

•F8266 

I 

•89924 

•89613 

•89304 

•88997 

•88692  j  '88389 

2 

•90050 

•89737 

•89428 

•89120 

•88815 

•88511 

3 

•90174 

•89862 

•89551 

•89243 

•88938 

•88634 

4 

•90299 

•89986 

•89675 

'89367 

•89061 

•88757 

5 

•90423 

•9OIIO 

•89799 

•89490 

•89184 

•88879 

o 

•90548 

•90234 

•89923 

•89614 

•89307 

•89002 

7 

•90073 

•90359 

•90047 

•?9737 

•89430 

•89124 

8 

•90798 

•90483 

•90171 

•89861 

•89553 

'89247 

9 

•90922 

•90607 

•90295 

•89984 

•89676 

'89369 

730 

•91047 

•90732 

•90418 

•90107 

•89799 

•89492 

i 

•91172 

•90856 

•90542 

•90231 

•89922 

•89615 

2 

•91296 

•90980 

•90666 

•90354 

•90045 

•89737 

3 

•91421 

•91104 

•90790 

•90478 

•90168 

•89860 

4 

•91546 

•91229 

•90914 

•90601 

•90291 

•89982 

5 

•91671 

•91353 

•91038 

•90725 

•90414 

•90105 

0 

•91795 

•9M77 

•91162 

•90848 

•90537 

•90228 

7 

•91920 

•91602 

•91285 

•90971 

•90660 

•90350 

8 

•92045    -91726 

•91409 

•91095 

•90784 

•90473 

9 

•92169   !   -91850     '91533 

•91218 

•90907 

•90594 

464 


Appendix. 


mm* 

15°. 

16°. 

17°. 

18°. 

19°. 

20°. 

740 

•92294 

•91975 

•91657 

•91342 

•91030 

•90718 

i 

•92419 

•92099 

•9I78I 

•9M65 

•9II53 

•90841 

2 

•92223 

•91905 

•91589 

•91276 

•90963 

3 

•92668 

•92347 

•92029 

•9I7I2 

•91399 

•91036 

4 

•92793 

•92472 

•92152 

•91836 

•91522 

•9I2O8 

5 

•92918 

•92596 

•92276 

•91959 

•91645 

•9I33I 

6 

•93043 

•92720 

•92400 

•92082 

•91768 

•9H54 

7 

•93167 

•92845 

•92524 

•92206 

•91891 

•91576 

8 

•93292 

•92969 

•92648 

•92329 

•92014 

•91699 

9 

•93417 

•93093 

•92772 

•92453 

•92137 

•9I82I 

750 

'93541 

•93217 

•92896 

•92576 

•92260 

•91944 

i 

•93666 

•93342 

•93O2O 

•92700 

•92383 

•92O67 

2 

•93791 

•93466 

•93H3 

•92823 

•92506 

•92189 

3 

•93916 

•93590 

•93267 

•92946 

•92629 

•92312 

4 

•91040 

•93715 

•93391 

•93070 

•92752 

•92434 

5 

•94165 

•93839 

•93515 

•93193 

•92875 

•92557 

6 

•94290 

•93963 

•93639 

•93317 

•92998 

•92679 

7 

•94414 

•94087 

•93763 

•93440 

•93I2I 

•92802 

8 

'94539 

•94212 

•93887 

•93564 

•93244 

•92925 

9 

•94664 

•94336 

•94OIO 

•93687 

•93367 

•93047 

760 

•9*789 

•94460 

•94134 

•938U 

•93489 

•93170 

i 

•94913 

•94585 

•94258 

'93934 

•93613 

•93292 

2 

•95038 

•94709 

•94382 

•94057 

•93736 

•93415 

3 

•95163 

•94833 

•94506 

•94181 

•93859 

•93538 

4 

•95288 

'94957 

•94630 

•94304 

•93982 

•93660 

5 

'954!2 

•95082 

'94754 

•94428 

•94105 

•93783 

6 

"95537 

•95206 

•94877 

•94551 

•94228 

•93905 

7 

•95662 

•95330 

•95001 

•94675 

•94351 

•94028 

8 

•95786 

'95455 

•95125 

•94798 

'94474 

•94I5I 

9 

•959H 

'95579 

•95249 

•94921 

'94597 

•94273 

770 

i 

•96036 
•96161 

•95703 
•95827 

'95373 
'95497 

•95045 
•95168 

•94720 
•94843 

•94396 
'945  '8 

2 

•96285 

•95952 

•95621 

•95292 

•94966 

•94641 

3 

•96410 

•96076 

'95744 

•95415 

•95089 

•94764 

4 

•96535 

•96200 

•95868 

'95539 

•95212 

•94886 

•96659 

•96325 

•95992 

•95662 

'95335 

•95009 

6 

•96784 

•96449 

•96116 

•95785 

•95458 

•95!3i 

7 

•96909 

•96573 

•96240 

•95909 

•9558i 

•95254 

8 

•97034 

•96698 

•96364 

•96032 

•95704 

•95376 

9 

•97158 

•96822 

•96488 

•96156 

•95827 

'95499 

78o 

•97283 

•96946 

•96611 

•96279 

•95950 

•95622 

Factors  for  reducing  Gaseous  Volumes  to  N.T.P.    465 


mm. 

21°. 

22°. 

23°- 

mm. 

21°. 

22°. 

23°. 

720 

•87972 

•87668 

•87371 

751 

•91759 

•91442 

•9H33 

•88094 

•87790 

•87493 

2 

•91882 

•91564 

•91254 

2 

•88216 

•879II 

•87614 

3 

•92004 

•91686 

3 

•88339 

•88033 

•87735 

4 

•92126 

•91808 

'9  '497 

4 

•88460 

•88155 

•87857 

5 

•92248 

•91929 

•91618 

5 

•88583 

•88277 

•87978 

6 

•92370 

•92051 

•91740 

6 

•88705 

•88398 

•88099 

7 

•92492 

•92173 

•91861 

7 

•88827 

•88520 

•88221 

8 

•92615 

•92295 

•91982 

8 

•88949 

•88642 

'88342 

9 

•92737 

•92416 

•92104 

9 

•89071 

'88764 

•88463 

760 

•92859 

•92538 

•92225 

73° 

•89193 

•88885 

•88585 

i 

•92981 

•92660 

•92346 

i 

•89316 

•89007 

•88706 

2 

•93103 

•92782 

•92468 

2 

3 

•89438 
•89560 

•89129 

•89251 

•88827 
•88949 

3 
4 

•93226 
•93348 

•92904 
•93025 

•92589 
•92711 

4 

•89682 

•89372 

'890,0 

5 

•93470 

•93147 

•92832 

5 

"89804 

•89494 

•89I9I 

6 

•93592 

•93269 

•92953 

6 

•89927 

•80616 

•89313 

7 

•937H 

•93391 

•93075 

7 

•90049 

•89738 

•89434 

8 

•93836 

•93512 

•93196 

8 

•90171 

'89860 

•89555 

9 

'93959 

•93634 

•93317 

9 

•90293 

•89981 

•89677 

770 

•94O8l 

•93756 

'93439 

740 

•90103 

•89798 

i 

•94203 

•93878 

•9356o 

i 

•^ 

•90225 

•89920 

2 

•94325 

'93999 

•93681 

2 

•90660 

•90347 

•90041 

3 

'94447 

1  94  1  2  1 

•93803 

3 

•90782 

•90468 

•90162 

4 

•945/0 

•94243 

•93924 

4 

•90904 

•9O59O 

•90284 

5 

•94692 

•94365 

•94045 

5 

•91026 

•90712 

•90405 

6 

•94814 

•94486 

•94167 

6 

•9II48 

•90834 

•90526 

7 

•94936 

•94608 

•94288 

7 

•9I27I 

•90955 

•90648 

8 

•95059 

•94730 

•94409 

8 

•91393 

•91077 

•90/69 

9 

•95180 

•94852 

•94531 

9 

•9II99 

•90890 

780 

•95303 

•94973 

•94652 

750 

•91637 

•9I32I 

•9IOI2 

466 


Appendix. 
LOGARITHMS.* 


0 

1 

2 

3 

4 

5 

6 

7 

8 

9 

123 

456 

789 

10 

cooo 

0043 

0086 

0128 

0170 

0212 

0253 

0294 

0334 

0374 

4  8  12 

17  21  25 

29  33  37 

II 

0414 

0453 

0492 

053  i 

0569 

0607 

0645 

0682 

0719 

°755 

4  8  ii 

15  19  23 

26  30  34 

.12 

0792 

0828 

0864 

0899 

°934 

0969 

1004 

1038 

1072 

1106 

3  7  10 

14  17  21 

24  28  31 

13 

H39 

H73 

1206 

1239 

1271 

1303 

1335 

1367 

1399 

143° 

3  6  10 

13  l6  19 

23  26  29 

14 

1461 

1492 

1523 

1553 

1584 

1614 

1644 

1673 

1703 

1732 

3  6  9 

12  15  18 

21  24  27 

15 

1761 

1790 

ibi8 

1847 

1875 

1903 

i93i 

1959 

1987 

2014 

368 

II  14  17 

20  22  25 

16 

2041 

2068 

2095 

2122 

2148 

2175 

2201 

2227 

2253 

2279 

3  5  8 

ii  13  16 

l8  21  24 

17 

2304 

2330 

2355 

2380 

2405 

2430 

2455 

2480 

2504 

2529 

257 

1O  12  15 

17  2O  22 

18 

2553 

2577 

2601 

2625 

2648 

2672 

2695 

2718 

2742 

2765 

257 

9  12  14 

16  19  21 

19 

2788 

2810 

2833 

2856 

2878 

2900 

2923 

2945 

2967 

2989 

247 

9  ii  13 

16  18  20 

20 

3010 

3032 

3054 

3075 

3096 

3Il8 

3139 

3160 

3181 

3201 

2  4  6 

8  ii  13 

15  17  19 

21 

3222 

3243 

3263 

3284 

3304 

3324 

3345 

3365 

3385 

3404 

246 

8  IO  12 

14  16  18 

22 

3424 

3444 

3464 

3483 

35°2 

3522 

354i 

3560 

3579 

3598 

246 

8  10  12 

H  15  17 

23 

3617 

3636 

3655 

3674 

3692 

37" 

3729 

3747 

3766 

3784 

246 

7  9  ii 

!3  15  17 

24 

3802 

3820 

3838 

3856 

3874 

3892 

3909 

3927 

3945 

3962 

245 

7  9  ii 

12  14  16 

25 

3979 

3997 

4014 

4°3i 

4048 

4065 

4082 

4099 

4116 

4i33 

235 

7  9  10 

12  14  15 

26 

4150 

4166 

4183 

4200 

4216 

4232 

4249 

4265 

4281 

4298 

235 

7  8  10 

ii  i3  i.5 

27 

4314 

4330 

4346 

4362 

4378 

4393 

4409 

4425 

4440 

4456 

235 

689 

ii  13  14 

28 

4472 

4487 

4502 

45i8 

4533 

4548 

4564 

4579 

4594 

4609 

235 

689 

II  12  14 

29 

4624 

4639 

4654 

4669 

4683 

4698 

4713 

4728 

4742 

4757 

1  3  4 

679 

10  12  13 

30 

477i 

4786 

4800 

4814 

4829 

4843 

4857 

4871 

4886 

4900 

i  3  4 

679 

10  ii  13 

31 

4914 

4928 

4942 

4955 

4969 

4983 

4997 

5011 

5024 

5038 

1  3  4 

678 

IO  II  12 

32 

505i 

5065 

5°79 

5092 

5io5 

5H9 

5132 

5i45 

5159 

5172 

1  3  4 

5  7  8 

9  n  12 

33 

5185 

5i98 

5211 

5224 

5237 

5250 

5263 

5276 

5289 

5302 

1  3  4 

568 

9  10  12 

34 

53i5 

5328 

5340 

5353 

5366 

5378 

539i 

5403 

54i6 

5428 

i  3  4 

568 

9  10  ii 

35 

5441 

5453 

5465 

5478 

549° 

5502 

55H 

5527 

5539 

555i 

I  2   4 

5  6  7 

9  10  ii 

36 

5563 

5575 

5587 

5599 

5611 

5623 

5635 

5647 

5658 

5670 

124 

5  6  7 

8  10  ii 

57 

5682 

5694 

5705 

57i7 

5729 

574° 

5752 

5763 

5775 

5786 

I  2   3 

5  6  7 

8  9  10 

38 

5798 

5809 

5821 

5832 

5843 

5855 

5866 

5877 

5888 

5899 

I  2   3 

5  6  7 

8  9  10 

39 

59H 

5922 

5933 

5944 

5955 

5966 

5977 

5988 

5999 

6010 

I  2   3 

457 

8  9  10 

40 

6021 

6031 

6042 

6053 

6064 

6075 

6085 

6096 

6107 

6117 

I  2   3 

4  5  6 

8  9  10 

41 

6128 

6138 

6149 

6160 

6170 

6180 

6191 

6201 

6212 

6222 

I  2   3 

4  5  6 

789 

42 

6232 

6243 

6253 

6263 

6274 

6284 

6294 

6304 

6314 

6325 

I  2   3 

4  5  6 

789 

43 

6335 

6345 

6355 

6365 

6375 

6385 

6395 

6405 

64i5 

6425 

I  2   3 

4  5  6 

789 

44 

6435 

6444 

6454 

6464 

6474 

6484 

6493 

6503 

6513 

6^22 

I  2   3 

4  5  6 

789 

45 

6532 

6542 

655i 

6561 

6571 

6580 

6590 

6599 

6609 

66l8 

I  2   3 

4  5  6 

789 

46 

6628 

6637 

6646 

6656 

6665 

6675  6684 

6693 

6702 

6712 

I  2   3 

4  5  6 

7  7  8 

47 

6721 

6730  6739 

6749 

6758 

6767 

6776 

6785 

6794 

6803 

I  2   3 

455 

678 

48 

6812 

6821 

6830 

6839 

6848 

6857 

6866 

6875 

6884 

6893 

I  2   3 

445 

678 

49 

6902 

6911 

6920 

6928 

6937 

6946 

6955 

6964 

6972 

6981 

I  2   3 

445 

678 

50 

6990 

6998 

7007 

7016 

7024 

7033 

7042 

7050 

7059 

7067 

I  2   3 

345 

678 

51 

7076 

7084 

7°93 

7101 

7110 

7118 

7126 

7135 

7143 

7152 

I  2   3 

345 

678 

52 

7160 

7168 

7177 

7185 

7193 

7202 

7210 

7218 

7226 

7235 

122 

345 

677 

53 

7243 

7251 

7259 

7267 

7275 

7284 

7292 

7300 

7308 

73l6 

122 

345 

667 

54 

7324 

7332 

7340 

7348 

7356 

7364 

7372 

7380 

7388 

7396 

122 

345 

667 

*  The  following  Tables  of  Logarithms  and  Antilogarithms  are  reprinted  from  "  Mathematical 
Tables  for  Ready  Reference, "  by  permission  of  Mr.  Castle. 


Logarithms. 


467 


0 

1 

2 

3 

4 

5 

6 

7 

8 

9 

123 

456 

789 

55 

7404 

7412 

74'9 

7427 

7435 

7443 

745' 

7459 

7466 

7474 

122 

345 

5  6  7 

56 

57 

7482 

7559 

7490 
7566 

7497 
7574 

7505 
7582 

75J3 
7589 

7520 
7597 

7528 
7604 

7612 

7543 
7619 

7627 

I      2 

I       2 

345 
345 

5  6  7 
5  6  7 

58 

7634 

7642 

7649 

7657 

7664 

7672 

7679 

7686  7694 

7701 

I       2 

344 

5  6  7 

7709 

7723 

7731 

7738 

7745 

7752 

7760  7767 

7774 

I       2 

344 

5  6  7 

60 

7782 

7789 

7796 

7803 

7810 

7818 

7825 

7832 

7839 

7846 

I       2 

344 

566 

6i 

7853 

7860 

7868 

7875 

7882 

7889 

7896 

7903 

7910 

7917 

I      2 

344 

566 

62 

7924 

793  z 

7938 

7945 

7952 

7959 

7966 

7973  j  798o 

7987 

i     2:3  3  4 

566 

63 

7993 

8000 

8007 

8014 

8021 

8028 

8035 

$041  8048 

8055 

i     2|3  3  4 

556 

64 

8062 

8069 

8075 

8082 

8089 

8096 

8102 

Siogi  8116 

8122 

i    2:3  3  4  5  5  6 

65 

8129 

8136 

8142 

8149 

8156 

8162 

8169 

8176 

8182 

8189 

i    2334 

5  5  6 

66 

8195 

5  202 

8209 

8215 

8222 

8228 

8235 

8241 

8248 

8254 

i    2334 

5  5  6 

67 

8261 

8267  1  8274 

8280 

8287 

8293 

8299 

830618312  8319 

2334 

5  5  6 

68 

8325 

8331  833818344 

8351 

8357 

8363 

837°  8376 

8382 

2<3  3  4 

4  5  6 

69  8388 

8395 

8401  8407 

8414 

8420 

8426 

843218-39 

8445 

2234 

4  5  6 

70 

845i 

8457 

8463 

8470 

8476 

8482 

8488 

8494 

8500 

8506 

2 

234 

4  5  6 

71 

8513 

8519 

8525 

8531 

8537 

8543 

8549 

8555 

8561 

8567 

2 

234 

455 

72 

8573 

8579 

8585 

8591 

8597 

8603 

8609 

8615 

8621 

8627 

2:2  3  4 

455 

73 

8633 

8639  1  8645 

8651 

8657 

8663 

8669 

5675 

8681 

8686 

2;2  3  4 

455 

74 

8692 

8698 

8704 

8710 

8716 

8722 

8727 

8733 

8739 

8745 

2:2  3  4 

455 

75 

8751 

8756 

8762 

8768 

8774 

8779 

8785 

8791 

8797 

8802 

2.2  3  3 

455 

76 

8808 

8814 

8820 

8825 

8831 

8837 

8842 

8848 

8854 

8859 

2:2  3  3 

455 

77 

8865 

8871 

8876 

8882 

8887 

8893 

8899 

8904 

8910 

8915 

2|2  3  3 

445 

78 

8921 

8927 

8932 

8938 

8943 

8949 

8954 

8960 

8965 

8971 

2  1  2  3  3!  4  4  5 

79 

8976 

8982 

8987 

8993 

8998 

9004 

9009 

9015 

9020 

9025 

2  2  3  3|  4  4  5 

80 

9031 

9036 

9042 

9047 

9053 

9058 

9063 

9069 

9074 

9079 

2 

233 

i  4  4  5 

81 

9085 

9090 

9096 

9101 

9106 

9112 

9117 

9122 

9128 

9*33 

2 

233 

445 

82 

9138 

9143 

9149 

9154 

9159 

9165 

9170 

9i75 

9180 

9186 

2233 

445 

83 

9191 

9196 

9201 

9206 

9212 

9217 

9222 

9227 

9232 

9238 

2233 

445 

84 

9243 

9248 

9253 

9258 

9263 

9269 

9274 

9279 

9284 

9289 

233 

445 

85 

9294 

9299 

9304 

9309 

93*5 

9320 

9325 

9330 

9335 

9340 

I      2 

233 

445 

86 

9345 

9350 

9355 

9360 

9365 

9370 

9375 

9380 

9385 

939° 

I      2 

233 

445 

87 

9395 

9400 

9405 

9410 

94i5 

9420 

9425 

9430 

9435 

9440 

0      I 

223 

344 

88 

9445 

945° 

9455 

9460 

9465 

9469 

9474 

9479 

9484 

9489 

0      I 

223 

344 

89^ 

9494 

9499 

9504 

9509 

0513 

9523 

9528 

9533 

9538 

O      I 

223 

344 

90 

9542 

9547 

9552 

9557 

9562 

9566 

957i 

9576 

9586 

0      I 

223 

344 

91 

9590 

9595 

9600 

9605 

9609 

9614 

9619 

9624 

9628 

9633 

O      I 

223 

344 

92 

9638 

9643 

9647  i  9652 

9657 

9661 

9666 

9671 

9675 

9680 

0      I 

223 

344 

93 

9685 

9689 

9694 

9699 

9703 

9708 

9713 

9717 

9722 

9727 

0      I 

223 

344 

94 

973i 

973^ 

974i 

9745 

9750 

9754 

9759 

9763 

9768 

9773 

O      I 

223 

344 

95 

9777 

9782 

9786 

9791 

9795 

9800 

9805 

9809 

9814 

9818 

0      I 

223 

344 

96 

9823 

9827 

9832 

9836 

9841 

9845 

9850 

9854 

9859 

9863 

O      I 

223 

344 

97 

9868 

9872 

9877 

9881 

9886 

9890 

9894 

9899 

9903 

9908 

0      I 

223 

344 

98 

9912 

9917 

9921 

9926 

993° 

9934 

9939 

9943 

9948 

9952 

0      I 

223344 

99 

9956 

9961 

9965 

9969 

9974 

9978 

9983 

9987 

9991 

9996 

O      I 

223334 

468 


Appendix. 


ANTILOGARITHMS. 


0 

1 

2 

3 

4 

5 

6 

7 

8 

9 

123 

456 

789 

•00 

IOOD 

1002 

1005 

1007 

1009 

1012 

1014 

1016 

1019 

1021 

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2    3    3 

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235° 

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233 

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2630 

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4    5    6 

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2704 

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2723 

2729 

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4    5    6 

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2767 

2773 

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2805 

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2864 

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334 

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2897 
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2904 
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345 

5  6  7 

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4710 

4721 

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4753 

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4808 

4819 

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4920 

4932 

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5916 

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5943 

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5970 

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6761 

6776 

6792 

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6823 

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6918 

6934 

6950 

6966 

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7015 

7031 

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7079 

7096 

7112 

7129 

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7178 

7194 

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12  13  15 

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7244 

7261 

7278 

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7328 

7345 

7362 

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12  13  15 

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7413 

7430 

7447 

7464 

7482 

7499 

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12  14  16 

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7603 

7621 

7638 

7656 

7674 

7691 

7709 

7727 

7745 

4  5 

7  9  ii 

12  14  16 

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7762 

7780 

7798 

7816 

7834 

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7870 

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7907 

7925 

4  5 

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13  14  16 

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7980  7998 

8017 

8035 

8054 

8072 

8091 

8110 

2  4  6 

7  9  ii 

13  15  17 

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8147 

8166  8185 

8204 

8222 

8241 

8260 

8279 

8299 

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8318 

8337 

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84M 

8433 

8453 

8472 

8492 

246 

8  10  12 

14  15  17 

"93 

8511 

8531 

8551 

8570 

8590 

8610 

8630 

8650 

8670 

8690 

2  4  6 

8  10  12 

14  16  18 

'94 

8710 

8730 

875° 

8770 

8790 

8810 

8831 

8851 

8872 

8892 

246 

8  10  12 

14  16  18 

'95 

8913 

8933 

8954 

8974 

8995 

9016 

9036 

9057 

9078 

9099 

2  4  6 

8  10  12 

IS  17  19 

•96 

9120 

9141 

9162 

9183 

9204 

9226 

9247 

9268 

9290 

9311 

246 

8  ii  13 

15  17  19 

'97 

9333 

9354 

9376 

9397 

9419  9441 

9462 

9484 

9506 

9528 

247 

9  ii  13 

15  17  20 

•98 

9550 

9572 

9594 

9616 

9638 

9661 

9683 

9705 

9727 

975° 

247 

9  i1  13 

16  18  20 

'99 

9772 

9795 

9817 

9840 

9863 

9886 

9908 

9954 

9977 

257 

9  ii  14 

16  18  20 

INDEX 


ABSORPTION   pipettes,  gas   analysis, 

393 

,  mode  of  filling,  394 

Acid,  acetic,  156 

,  boric,  169 

,  carbonic,  153 

,   ,    gasometric    estimation, 

368-396 

, ,  gravimetric  estimation, 259 

, ,  volumetric  estimation,  364 

,  citric,  158 

• ,  cyanic,  164 

,  formic,  155 

,  hydriodic,  128 

,  hydrobromic,  125 

,  hydrochloric,  122 

, ,  gravimetric  estimation,  259 

,  ,  standard  solution,  323 

— ,  hydrocyanic,  159 

,  hydrofluoric,  133 

,  hydrofluosilicic,  135 

,  hypochlorous,  129 

,  hypophosphorous,  151 

— ,  metaphosphoric,  150 

— ,  metastannic,  104 

— ,  nitric,  144 

, .gravimetric estimation,  263 

,  ,  volumetric  estimation,  363 

,  nitrous,  147 

,  orthophosphoric,  150 

,  osmic,  109 

,  oxalic,  155 

, ,  standard  solution,  378 

— ,  phosphoric,  63,  149 

, ,  gravimetric  estimation,  264 

, ,  volumetric  estimation,  369 

,  phosphorous,  151 

— ,  pyrophosphoric,  150 

— ,  radicals,  15 

,  silicic,  166 

,  sulphuric,  139 

— , ,  gravimetric  estimation,  258 

, ,  standard  solution,  319 

— ,  tartaric,  157 
,  thiocyanic,  165 

— ,  thiosulphuric,  142 
Acetates,  156 
Acidimetry,  316 


Acids,  analytical  classification  of,  171 

Acids  and  alkalies,  12 

Agate,  1 66 

Air-oven,  198 

Alkalies,    estimation   of   in   silicates, 

287 

, by  electrolytic  method,  290 

Alkalimetry,  316 

Alloy,  silver  and  copper,  electrolytic 

analysis,  297 
, , ,  gravimetric  analysis, 

267 

,  tin  and  lead,  gravimetric  analy- 
sis, 269 
,  zinc,  copper,  nickel,  electrolytic 

analysis,  297 

— , , ,  gravimetric  analysis, 

273 

,  zinc,  copper,  tin,  274 

Aluminium,  gravimetric  estimation  in 

dolomite,  277 

— ,  —          —  potash  alum,  222 

reactions,  37 

Alums,  37 

Ammonia,  volumetric  estimation,  327 

Ammonium,   gravimetric   estimation, 

250 

— ,  phosphomolybdate,  63 

— ,  reactions,  19 

*- ,   thiocyanate,  deci-normal   solu- 
tion, 365 

Analytical  classification,  13 
Analytical  groups,  16 
Analytical  tables — 

General  table,  i85A 

Group  I.,  115 

II.,  Division  i,  88 

II.,  Division  2,  108 

IIIA,  49 

II IB,  62 

III.  \pho3phate  table),  1858 

IV.,  35 

• V.,  25 

Antimoniuretted  hydrogen,  101 
Antimony,  gravimetric  estimation,  255 
, ,  separation   from  arsenic, 

284 

—  reactions,  97 


472 


Index, 


Antimony,  volumetric  estimation,  352 

Aqueous  vapour,  tension  of,  461 

Arsenic,  gravimetric  estimation  as  sul- 
phide, 253 

,   - magnesium    pyro- 

arsenate,  254 

reactions,  89 

— ,  volumetric  estimation,  352 

Aspirator,  204 

Atomic  weights,  table  of,  453 

BALANCE,  190 

Barium,  gravimetric  estimation,  258 

Barium  hydroxide,  standard  solution, 

379 

Barium  reactions,  30 
— ,  spectrum  of,  34 

Baryta-water,  379 

Beryl,  70 

Beryllium,  71 

Bismuth,  reactions,  81 

Bleaching  powder,  130 

,  available  chlorine  in,  353 

,  valuation  of  by  arsenious 

oxide,  354 

,  valuation  by  means  of  gas- 
volumeter,  412 

Blowpipe  flame,  9 

Blowpipe  reactions,  exercises,  9 

Boiling-point  determination,  436 
—  method,  448 

Borates,  169 

Borax,  169 

Borax  beads,  n 

Boron,  169 

Boron  fluoride,  135,  170 

Bromates,  130 
— ,  volumetric  estimation,  358 

Bromides,  125 

Bromides,    'iodides,    and     chlorides, 
detection  in  solution  together,  128 

Bromine,  124 

,   gravimetric    estimates  in  bro- 
mides, 259 

,    estimation    in    organic     com- 
pounds, 430 

Bronze  coinage,  analysis  of,  274 

Burettes,  308 

,  calibration  of,  309 

,  gas,  385 

CADMIUM, gravimetric  estimation,  257 

,  —        —  in  zinc-blende,  282 

,  reactions,  83 

— ,  volumetric  estimation,  364 
Caesium,  26 
Calcium,    gravimetric   estimation    in 

calcium  carbonate,  231 

, dolomite,  277 

, silicates,  287 


Calcium  reactions,  32 
— ,  volumetric  estimation,  340,  363 

Calibration  of  burettes,  309 
gas-burettes,  386 

litre  flasks,  305 

— —  pipettes,  307 

Carbon,  153 

,  estimation  in  organic  com- 
pounds, 414 

Carbon  dioxide,  estimation  by  means 
of  gas-volumetre,  412 

,  gravimetric  estimation,  259 

,  indirect  volumetric  estima- 
tion, 364 

in     air    (Pettenkofer's 

method),  378 
in  other  gas  mixtures,  399 

Carbon  monoxide,  estimation  of,  401 

Carbonates,  153 

Cerium,  16 

Chlorates,  130 

and  nitrates,  detection  together, 

J47 

,  separation  from  chlorides,  133 

— ,  volumetric  estimation,  358 

Chlorides,  121 

,  detection  in  presence  of  bromide 

or  iodide,  124 

Chlorine,  121 

,  gravimetric  estimation  in  chlo- 
rides, 259 

,  in  organic  com- 
pounds, 430 

,  volumetric  estimation  by  pre- 
cipitation, 360,  366 

, in  bleaching  powder, 

353 

Chrome  iron  ore,  39,  344 

Chromium,  gravimetric  estimation, 
229 

reactions,  39 

,  volumetric  estimation  in  chrome 

iron  ore,  344 

Chromyl  chloride,  123 

Citrates,  158 

Clark's  method  for  estimating  hard- 
ness of  water,  371 

Coal-gas,  analysis  of,  408 

Cobalt,  gravimetric  estimation,  247 

,  reactions,  58 

Cobaltamines,  59 

Cobalticyanides,  60 

Cobaltocyanides,  60 

Coefficients  of  expansion  of  water, 
table  of,  460 

Collection  of  gases  for  analysis,  390 

Combustion  analysis,  414 

Combustion  tube,  414 

Copper,  electrolytic  estimation,  294 


Index. 


473 


Copper,  gravimetric  estimation  in 
copper  sulphate,  (i)  as  oxide,  235  ; 
(2)  as  sulphide,  236  ;  (3)  as  cuprous 
thio-cyanate,  237 

f in  German  silver,  271 

\ silver  coinage,  267 

t bronze  coinage,  275 

i zinc-blende,  282 

'reactions,  83 

1   volumetric    estimation   (Mans- 
field process),  375 
Correction  of  gaseous  volumes,  387 

,  factors  for,  462 

Crucible  tongs,  manipulation  of,  227 

Cryoscopic  method,  443 

Crystallisation,  208 

Cupric  salts,  85 

Cuprous  chloride  for  gas  absorption, 

396 

hydride,  151 

salts,  86 

Cyanates,  164 
Cyanides,  double,  161 

—  of  Group  II IB.,  60 

— ,  reactions,  159 

,  single,  160 

Cyanogen,  159 

(  volumetric  estimation,  362 

DESICCATOR,  196 
Dolomite,  analysis  of,  276 
Double  cyanides,  161 

easily  decomposed,  162 

difficultly  decomposed,  163 

Double  gas-pipette,  mode  of  filling, 

394 

Double  salts,  213 
Drying  and  weighing  a  filter,  193 

ELECTROLYTIC  methods  of  analysis, 
292 

Etching  glass,  134 

Evaporating  to  dry  ness,  4 

Evaporation,  4 

Expansion  of  water,  table  of  coeffi- 
cients, 460 

Explosion-pipettes,  404,  407 

FACTORS  for  correction  of  gaseous 

volumes,  462 
Feeling's  solution,  86 
Felspar,  167 
Ferricyanides,  164 
Ferrocyanides,  163 
Filter,  incineration  of,  207 
Filter  ash,  estimation  of,  206 
Filter-paper,  quantitative,  193 
Filter-pump,  212 
Filtration,  i 
Flame,  oxidising  and  reducing,  9 


Fleitmann's  test,  97 

Flint,  166 

Floats,  311 

Fluid  magnesia,  153 

Fluorides,  133 

Fluorine,  133 

Formates,  155 

Freezing-point  method,  442 

Fulminating  gold,  106 

Furnace  gases,  sulphur  dioxide  in,  381 

Fusion,  5 

Fusion  mixture,  49 

Fusion  with  borax,  exercises,  10 

GAS  analysis,  377 
Gas-burettes,  385 
Gas- volumeter,  411 
Gases,  absorption  analysis,  396,  398 
,  correction  of  volumes  for  tem- 
perature and  pressure,  387 
,   estimation  by  absorption  and 

titration,  378 

, measurement,  384 

,  measurement  of,  384 

Geissler's  potash  bulbs,  417 

General  reagents,  17 

General  table  for  the  separation  of 

the  metals,  i85A 
German  silver,  electrolytic  analysis  of, 

297 

,  gravimetric  analysis  of,  270 

Gold  reactions,  105 
Graduated  cylinders,  306 
Graduated  flasks,  303 
Gravimetric  methods  of  analysis,  189 
Gravimetric  separations — 

Antimony  and  arsenic,  284 

Copper  and  cadmium,  282 

silver,  267 

nickel,  271 

zinc,  271 

Iron  and  aluminium,  278 

manganese,  283 

zinc,  271 

Magnesium  and  calcium,  277 

Manganese  and  zinc,  283 

Nickel  and  zinc,  272 

Tin  and  copper,  272 

Tin  and  lead,  269 
Group  separations,  116 
Group  V. ,  detection  of  metals  of,  25 

IV. ,  separation  of  metals  of,  35 

IIIA,  „  ,,         49 

1KB,  ,,  ,,62 

II.,    Division   i,    separation    of 

metals  of,  88 

—  II.,    Division    2,    separation    of 

metals  of,  108 

I.,  separation  of  metals  of,  115 

Group  reagents,  17 


474 


Index. 


HALIDES,  121 

Halogen  oxyacids,  129 

Halogens,  121 

Hardness  of  water,   Clark's  process, 

,  Hehner's  method,  328 

Hardness,  table  of,  459 
Hempel's  gas-burette.  385 
Hot- water  funnel,  210 
Hydrogen,     estimation    in     gaseous 
mixtures,  402,  404,  408 

, organic  compounds,  414 

Hypochlorites,  129 
Hypochlorous  acid,  129 
Hypophosphites,  151 


IGNITION,  8 
Indicators,  use  of,  315 

for  alkalimetry,  317 

Indium,  16 

Insoluble  substances,  treatment  of,  184 
Instruments  for  measuring  liquids,  303 
lodates,  130 

,  volumetric  estimation  of,  358 

Iodides,  129 

,     detection     in    solution     with 

bromides,  127 
Iodine,  126 

,  deci-norma]  solution,  348 

,  estimations  by  means  of,  352 

,  separation  from   bromine    and 

chlorine,  129 

,  volumetric  reactions  with,  346 

Ions,  292 

Iridium,  109 

Iron,    gravimetric    estimation    in    a 

ferrous  salt,  230 

, German  silver,  272 

, dolomite,  278 

, zinc-blende,  282 

— ,  magnetic  oxide,  44 

reactions,  44 

,  volumetric  estimation  of,  in  ores, 

335-  343 


LEAD,  gravimetric  estimation  in  lead 
acetate,  (i)  as  oxide,  240;  (2)35 
sulphate,  242 

( bronze,  274 

>, solder,  270 

, zinc-blende,  281 

reactions,  79 

Lithium  reactions,  26 
— ,  spectrum  of,  28 

Litmus,  12,  316 

Litre  flasks,  calibration  of,  305 

Logarithms,  table  of,  466 


MAGNESIA  mixture,  456 

Magnesium,  gravimetric  estimation  in 
magnesium  sulphate,  233 

, dolomite,  278 

, silicates,  287 

,  pyro-arsenate,  254 

—  reactions,  23 

Manganese  dioxide,  valuation  of,  337, 
356 

,  gravimetric  estimation  in  potas- 
sium permanganate,  245 

, zinc-blende,  283 

reactions,  51 

Mansfield  process  for  copper,  375 

Marsh-gas,  estimation  of,  406,  408 

Marsh's  test,  94 

Melting-point  determination,  438 

Meniscus,  310 

Mercuric  compounds,  reactions,  76 

Mercurous  compounds,  78 

Mercury,  gravimetric  estimation,  257 

Mercury  reactions,  76 

Metals,'  15 

Metalloidal  elements,  16 

Meta- phosphates,  150 

Methane,  estimation  of,  406,  408 

Methyl  orange,  317 

Miscellaneous  physico-chemical  deter- 
minations, 434 

Molecular  weight  by  boiling-point,  448 

Molecular  weight  by  freezing-point, 442 

Molybdenum,  in 

NEGATIVE  radicals,  15 
Nessler's  solution,  20 
Neutralisation,  exercises,  n 
Nickel,  electrolytic  estimation  of,  297, 
299 

,  gravimetric  estimation,  247 

, in  German  silver,  272 

reactions,  56 

Niobium,  16 
Nitrates,  144 
and  chlorates,  detection  together, 

T47 
and  nitrites.detection  together,  149 

,  reduction  by  ferrous  salts,  145 

— ,  nascent  hydrogen,  146 

— ,  sulphurous  acid,  145 

Nitric  acid,  gravimetric  estimation  of, 

263 

— ,  volumetric  estimation,  363,  411 
Nitrites,  147 

— ,  distinction  from  nitrates,  148 
Nitrogen,     estimation     in     gaseous 

mixture,  408 

, organic  compounds,  424 

,  —  —  by  Dumas'   method, 

424 


Index. 


475 


Nitrogen,  estimation  in  organic  com- 
pounds by  Kjeldahl's  method,  429 

,  —  by  soda-lime  process, 

427 

Nitrometer,  410 

Nitroprusside  of  sodium,  139 

ORTHOPHOSPHATES,  150 

Osmium,  109 

Oxalates,  155 

Oxalic  acid,  standard  for  gas  analysis, 

378 
Oxygen,  estimation  of,  401 

FALLACIOUS  nitrate,  129 

Palladium,  109 

Palladium ised  asbestos,  402 

Perchlorates,  133 

Pettenkofer's  method  for  carbon 
dioxide,  378 

Phenol phthalein,  317 

Phosphites,  151 

Phosphates,  Group  III.,  63 

,  reactions  of,  63,  149 

Phosphoric  acid,  gravimetric  estima- 
tion of,  264 

,  removal  of,  in  Group  III.,  67 

,  separation  of,  69 . 

,  volumetric  estimation,  369 

Phosphorus,  149 

,  estimation  in  organic  com- 
pounds, 432 

Physico-chemical  determinations,  434 

Pipettes,  306 

,  calibration  of,  307 

Platinum  reactions,  105 

Positive  radicals,  15 

Potassioscope,  22,  33 

Potassium  cyanide,  standard  solution 
for  copper,  374 

dichromate,  deci-normal,  341 

,  analyses  by  means  of,  342 

,  gravimetric  estimation  of,  (i)as 

potassium    sulphate,    247 ;    (2)   as 
potassium  platinic  chloride,  249 

, in  silicates,  287 

Potassium  oleate,  standard  solution, 
372 

permanganate,  deci-normal  solu- 
tion, 332 
,  typical  analyses  by  means 

of,  335 

Potassium  reactions,  22 
Powdering  and  sampling,  265 
Precipitation,  6 

Preliminary  exercises,  qualitative,  i 
Preliminary  manipulations,  quantita- 
tive, 189 


Preliminary  tests  in  qualitative  ana- 
lysis, 178 

Pure  salts,  preparation  of,  208 
Purple  of  Cassius,  103,  107 
Pyrogallol,  397 
Pyrophosphates,  150 

QUANTITATIVE  analysis,  189 
Quartz,  166 

RAOULT'S  method,  443 
Rare  metals  of  Group  I.,  118 

II.,  108 

III.,  70 

Reactions  of  the  metals  of  Group  I. , 
112 

II.,  Division  i,  76 

• II.,  Division  2,  89 

-IIIA.,37 

HIB.,51 

IV.,  30 

V.,  19 

Reactions,  wet  and  dry,  13 
Reagents,  14 

,  group  or  general,  17 

,  preparation  of,  454 

Reinsch's  test,  92 

Results  of  a  qualitative  analysis,  186 

Rhodium,  16,  108 

Rose's  crucible,  237 

Rubidium,  26 

Ruthenium,  16,  108 

SAMPLING  and  powdering,  265 

Sand,  1 66 

Scandium,  16 

Scheele's  green,  92 

Schiff's  burette,  426 

Schrotter's     apparatus     for     carbon 

dioxide,  260 
Selenium,  no 
Separation  of  phosphoric  acid,  69 

the  metals  of  Group  I.,  115 

II.,  Division  i,  88 

II.,  Division  2,  108 

IIIA.,49 

IIlB.,  62 

IV.,  35 

Silica,  166 

,  estimation  of,  in  silicates,  285 

Silicates,  166 

,  gravimetric  analysis  of,  285 

,  treatment  of  insoluble,  168 

Silicic   acid,   gravimetric  estimation, 

264,  285 

Silicofluorides,  135 
Silicon,  166 
Silicon  fluoride,  134 
Silver  coinage,  267 


476 


Index, 


Silver  coinage,  electrolytic  analysis  of, 

297 
Silver,  gravimetric  estimation  in  silver 

nitrate,  238 

coinage,  267 

Silver  reactions,  112 

nitrate,  deci-nortnal,  360 

volumetric  estimation  of,  365,  363 

Single  cyanides,  160 

Soap  solution,  standard,  372 

Soda-ash,  volumetric  estimation  of ,  324 

Sodium  nitroprusside,  139 

chloride,  deci-normal  solution, 

364 
,    gravimetric    estimation    in    a 

silicate,  287 
,    gravimetric    separation     from 

potassium,  288 

reactions,  21 

sulphide,  standard  solution,  370 

—  thiosulphate,  deci-normal,  348 
Solder,  gravimetric  analysis  of,  269 
Soluble  glass,  167 
S'olubilities,  table  of,  177 
Solvent,  3 
Solution,  3 

Specific-gravity  determinations,  434 
Spectra  of  alkali  metals,  28 

alkaline  earth  metals,  34 

Spectroscope,  27 
Spontaneous  evaporation,  4 
Standard  solutions,  312 
Stannic  compounds,  103 
Stannous  compounds,  102 
Steam-oven,  195 
Strontium  reactions,  31 

— ,  spectrum  of,  34 
Sulphates,  139 
Sulphides,  137 
Sulphites,  140 
Sulphur,  137 

,    estimation    in    organic    com- 
pounds, 432 

,  estimation  in  zinc-blende,  280 

dioxide,  estimation  by  means  of 

iodine,  353 

in  furnace  gases,  381 

Sulphuretted  hydrogen,  137 
Sulphuric  acid,  gravimetric  estimation 

of,  258 

,  standard  solution,  319 

Systematic  detection  of  the  acids,  171 
separation  of  the  groups,  116 


TARTAR  emetic,  98 

Tartrates,  157 

Tellurium,  no 

Tension  of  aqueous  vapour,  table  of, 

461 

Test-mixer,  306 
Thallium,  118 
Thiocyanates,  165 
Thiosulphates,  142 
Thorium,  16 
Tin,  gravimetric  estimation,  252 

, in  solder,  270 

, in  German  silver,  272 

, in  bronze  coinage,  274 

reactions,  101 

,    volumetric  estimation  of,  339, 

352 

Titanium,  73 
Tungsten,  119 

URANIUM,  74' 

acetate,  standard  solution,  367 

VANADIUM,  18 

Vapour-density  determination,  439 

Volumetric  methods  of  analysis,  300 

based  on  oxidation,  331 

precipitation,  360 

WATER,  hardness  of,  328,  371 
of   crystallisation,    direct    esti- 
mation, 203 

,  estimation  by  loss,  197 

Weighing,  189 
Weighing-bottle,  193 
Weights-,  191 

YTTERBIUM,  16 
Yttrium,  16 

ZINC,  electrolytic  estimation  of,  296 
,  gravimetric  estimation   in  zinc 

sulphate,  (i)  as  oxide,  243;  (2)  as 

sulphide,  244 

,  —          —  German  silver,  271 

, bronze,  275 

— ,  —        —  zinc-blende,  283 

—  reactions,  54 

— ,  volumetric  estimation,  371 
Zinc-blende,  analysis  of,  279 
Z.inci  carbonas,  55 
Zirconium,  72 


THE    END 


PRINTED    BY    WILLIAM    CLOWES    ANU   SONS,    LIMITED,    LONDON    AND    BECCLIS. 


OP   THF 

UNIVERSITY 


UNIVERSITY  OF  CALIFORNIA  LIBRARY 
BERKELEY 

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